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Full text of "General chemistry"







c ]\fa Qraw-3/ill Book & 7m 


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Plate of Spectra 









6 & 8 BOUVERIE ST., E. C. 





This book is something of an abridgment and much of a sim- 
plification of the author's earlier book entitled "Inorganic Chemis- 
try." It is intended for classes which have less time to devote 
to the subject than those for whom the former book was pre- 
pared. In it the author has preserved the general arrangement 
and spirit of the "Inorganic Chemistry," and has avoided a long 
more or less theoretical introduction, developing the subject as 
logically as possible from the descriptive and experimental side; 
each law and theory has been presented at the point which seems 
best fitted both to the student and to the subject. 

Perhaps the most marked difference from the former book is 
to be found in the treatment of the atomic theory. The author 
feels that the recent advancements have justified a more un- 
qualified support of this theory than was accorded it in the 
former work. By introducing an experimental conception of 
molar weights before taking up the chapter on atomic weights, 
the author feels that he has been able to develop the conception 
of atomic weights in a very simple manner from the experimental 
standpoint. The theory is then brought in as an explanation 
for the facts previously discussed. 

In the preparation of the present book, the author has gladly 
availed himself of all the suggestions which have come to him 
from others who have had experience with his former book. He 
feels very grateful to Dr. Paul V. Faragher for advice in some 
difficult places, and especially so to his wife, Stella G. Cady, 
for loyal assistance in preparing the manuscript and reading 
the proof. 


September, 1916. 







Properties of Substances Changes Natural Laws The Fun- 
damental Law of Chemistry Properties of a Substance and the 
Substance Physical and Chemical Properties Identification of 
Substances The Characterization of a Substance Physical States 
The Solid State The Liquid State The Gaseous State Solu- 
bility and Solution Mechanical Mixtures. 



Units Length Mass Weight Volume Energy Conservation 
of Weight and of Mass Density Specific Gravity. 


OXYGEN < . 12 

Discovery of Oxygen Combustion Role of the Air Com- 
bination and Decomposition Preparation of Oxygen Physical 
Properties Standard Conditions The Gas Laws Boyle's 
Absolute Temperature Other Physical Properties of Oxygen 
Liquid Oxygen Solid Oxygen Chemical Properties Quantita- 
tive Relations Elements Compounds Occurrence in Nature 
Preparation of Ozone Physical Properties of Ozone Chemical 
Properties Uses for Ozone Allotropy. 



Occurrence Preparation of Hydrogen Preparation from Acids 
Electrolytic Preparation Purification of the Hydrogen Physical 
Properties Diffusion Hydrogen and the Gas Laws Liquid 
Hydrogen Solid Hydrogen Absorption of Hydrogen by Metals 
Spectrum of Hydrogen Chemical Properties of Hydrogen 
Oxy-hydrogen Blowpipe Formation of Water from Hydrogen and 
Oxygen Compounds Catalysis. 






Physical Properties of Water Ice Properties of Ice Heat of 
Fusion Measurement of Heat Gaseous Water Humidity of the 
Air Heat of Vaporization of Water Effect of Pressure on the 
Melting-point of Ice The Law of Mobile Equilibrium Solutions 
Application of the Law of Mobile Equilibrium The Chemical 
Properties of Water The Reaction of Sodium and Water Th.e 
Composition of Water Molar Weight. 



Atomic Theory Molecular Theory Relation Between Fact and 
Theory Standard of Atomic Weights Symbols and Formulas 
Molecular Formulas of the Elements Chemical Equations. 



The Law of Multiple Proportions Preparation and Properties of 
Hydrogen Peroxide Oxidizing Action of Hydrogen Peroxide 
Reducing Action of Hydrogen Peroxide Determination of the 
Molecular Weight of Hydrogen Peroxide Composition and 
Formula of Hydrogen Peroxide Equation for the Preparation of 
Hydrogen Peroxide Explosive Properties of Hydrogen Peroxide 
Thermochemistry The Heat of Formation of Hydrogen Peroxide. 



Preparation of the Element Other Methods Technical Prep- 
aration Physical Properties Chemical Properties Chlorine 
Water Chlorine Hydrate The Naming of Chlorides Hydrogen 
Chloride Photochemical Action Action of Chlorine on Hydrogen 
Compounds Laboratory and Technical Preparation of Hydrogen 
Chloride First General Method for the Preparation of Acids 
Properties of Hydrogen Chloride Constant Boiling Hydrochloric 



Acids Bases Neutralization Reacting Ratio in Neutralization 
Acids, Bases, and Salts have two Sets of Properties Abnormal 
Lowering of the Freezing-point Electrolytes The Law of Fara- 
day Summary First Second Third The Theory of Elec- 
trolytic Dissociation Neutralization of Acids and Bases Heat 
of Neutralization Action of Acids Explanation Hydrogen 
Ion as an Oxidizing Agent. 





General Nomenclature of Acids and Salts The Preparation of 
the Oxygen Compounds of Chlorine Decomposition of Hypo- 
chlorites Hypochlorous Acid Properties of Hypochlorous Acid 
Chlorine Monoxide Chlorates Properties of Chlorates Prep- 
aration of Chloric Acid Separation of Salts Perchlorates 
Perchloric Acid Perchloric Anhydride Chlorine Dioxide and 
Chlorites Relative Oxidizing Power. 



General BROMINE Occurrence Preparation of the Element 
Physical Properties of Bromine Chemical Properties Hydrogen 
Bromide Properties of Hydrogen Bromide Oxygen Compounds 
of Bromine IODINE Preparation of the Element Physical Prop- 
erties of Iodine Law of Distribution Chemical Properties of 
Iodine Hydrogen Iodide Physical Properties Chemical Prop- 
erties Oxygen Compounds of Iodine Periodates and Periodic 
Acid Oxides of Iodine FLUORINE Occurrence Preparation 
Physical Properties Chemical Properties Hydrogen Fluoride 
Physical Properties Hydrolysis VALENCE Equivalent Weights. 



General Extraction of Sulfur Purification of Sulfur Crystalline 
Forms of Sulfur Transition Point Liquid Sulfur Amorphous 
Sulfur Molar Weight Chemical Properties of Sulfur Uses of 
Sulfur Hydrogen Sulfide Physical Properties Chemical Prop- 
erties Solution in Water Dissociation of Dibasic Acids Analyt- 
ical Reactions of Hydrogen Sulfide Hydrogen Sulfide as a Reduc- 
ing Agent Polysulfides Sulfur Dioxide or Sulfurous Anhydride 
Physical Properties Chemical Properties Reducing Action 
Bleaching Action Action on Organisms Sulfur Trioxide or 
Sulfuric Anhydride Action on Water Lead Chamber Process 
Concentration of the Acid Physical Properties Aqueous Sulfuric 
Acid Dissociation of Sulfuric Acid Applications of Sulfuric 
Acid Sulfates Identification Persulfuric Acid OTHER OXYGEN 
ACIDS OF SULFUR Hyposulfurous Acid Thiosulfuric Acid 
Halogen Compounds of Sulfur Sulfuryl Chloride Positive and 
Negative Valence Oxidation and Reduction. 



SELENIUM Properties Chemical Properties Hydrogen Selenide 
Selenium Dioxide and Selenious Acid Selenic Acid TELLURIUM 
Hydrogen Telluride Tellurium Dioxide Tellurous Acid. 





General Preparation of the Element Physical Properties 
Chemical Properties Hydrogen Compounds of Nitrogen Am- 
monia Occurrence Preparation Physical Properties Chem- 
ical Properties of Ammonia Hydrazine Hydronitric Acid or 
Hydrazoic Acid Air Composition of the Air Liquid Air Ex- 
periments Air is a Mixture Air and Life Oxides and Oxyacids 
of Nitrogen Nitric Acid Chemical Properties of Nitric Acid 
Aqua Regia Nitrates Nitro Compounds Nitrogen Pentoxide 
Nitric Oxide Properties of Nitric Oxide Nitrogen Peroxide 
Nitric Acid from the Air Nitrites Nitrous Anhydride Hypo- 
nitrites and Hyponitrous Acid Nitrous Oxide Nitrogen and Life. 


PHOSPHORUS .,.. . .. . 214 

General Occurrence Preparation of Phosphorus The Allotropic 
Modifications of Phosphorus Hydrogen Compounds of Phos- 
phorus Phosphonium Compounds Liquid Hydrogen Phosphide 
Halogen Compounds of Phosphorus Oxygen Compounds of 
Phosphorus Acids of Phosphorus Pyrophosphoric Acid Meta- 
phosphoric Acid Phosphorous Acid Phosphorus Sulfides 
Applications of Phosphorus. 



General Occurrence Diamond Graphite Amorphous Carbon 
Charcoal Relation Between the Allotropic Modifications Oxy- 
gen Compounds Carbon Dioxide Circulation of Carbon Photo- 
chemical Action Carbon Monoxide Properties Carbon Disul- 
fide Nitrogen Compounds of Carbon Cyanates Thyiocyan- 
ates The Hydrocarbons Crude Petroleum Methane Radicals 
Other Hydrocarbons Types of Carbon Compounds The 
Alcohols The Ethers The Aldehydes The Acids Acetic Acid 
The Esters Carbohydrates Some Explosives Unsaturated 
Hydrocarbons Acetylene Fuel Gases Coal Gas Water Gas- 
Producer Gas Blau Gas FLAMES Bunsen Burner. 


General Preparation Properties Silicon Dioxide Silicic Acid 
and Silicates Silicon Hydride Halogen Compounds Carborun- 




BORON 260 

General Boric Acid Other Compounds. 



Helium Neon Argon Krypton and Xenon Niton or Radium 



General Characteristic Physical Properties of Metals Charac- 
teristic Chemical Properties of Metals Distinction between 
Metals and Non-metals Classification of the Metals. 



SODIUM Occurrence Preparation Properties Sodium Hydride 
Sodium Chloride Sodium Bromide and Iodide Sodium Oxides 
Sodium Hydroxide Sodium Carbonate Solvay or Ammonia 
Soda Process Properties of Sodium Carbonate Uses of Sodium 
Carbonate Sodium Bicarbonate Sodium Nitrate Sodium 
Nitrite Sodium Sulfate Sodium Sulfite Sodium Sulfide 
Sodium Thiosulfate Oxyhalogen Compounds of Sodium Sodium 
Phosphates Sodium Tetraborate Sodium Silicate Sodium 
Cyanide Sodium Acetate Analytical Properties of Sodium 
POTASSIUM Occurrence Potassium Compounds and Plants Prep- 
aration of the Element Physical Properties Chemical Prop- 
erties Potassium Hydride Potassium Oxides Potassium Hy- 
droxide Potassium Chloride Potassium Bromide Potassium 
Iodide Potassium Chlorate Potassium Perchlorate Potassium 
Nitrate Potassium Carbonate Potassium Bicarbonate Potas- 
sium Cyanide Potassium Sulfate Analytical Properties of 
Hydroxide Ammonium Chloride Ammonium Nitrate Ammo- 
nium Sulfate The Sulfides of Ammonia Ammonium Carbonate 
Ammonium Oxalate Ammonium Thiocyanate Analytical Reac- 
tions of Ammonium Ammonium Amalgam LITHIUM Lithium 
Carbonate Lithium Chloride Analytical Properties of Lithium. 



General CALCIUM Occurrence Preparation of the Element 
Properties Calcium Hydride Calcium Oxide Calcium Hy- 
droxide Mortar Calcium Carbonate Calcium Chloride Cal- 
cium Hypochlorite and Bleaching Powder Calcium Bromide and 
Iodide Calcium Fluoride Calcium Nitrate Calcium Sulfate 
Calcium Sulfide Calcium Phosphate Calcium Carbide and 



Calcium Cyanamide Calcium Oxalate Calcium Silicate and 
Glass Analytical Properties of Calcium STRONTIUM Occur- 
rence The Compounds of Strontium BARIUM Occurrence 
Barium Oxide and Hydroxide Barium Dioxide Barium Carbon- 
ate Barium Sulfate Barium Sulfide Barium Chloride Barium 
Chlorate Barium Nitrate Barium Chromate Analytical Prop- 
erties of Barium RADIUM General Relationships. 



Magnesium Oxide and Hydroxide Magnesium Sulfate Mag- 
nesium Sulfide Phosphates of Magnesium Magnesium Nitride 
Analytical Reactions of Magnesium Water Softening ZINC 
General Occurrence Metallurgy Physical Properties Chem- 
ical Properties Galvanized Iron Zinc Oxide and Hydroxide 
Zinc Chloride Zinc Sulfate Zinc Carbonate Zinc Sulfide 
Analytical Properties of Zinc CADMIUM Cadmium Sulfide 
Analytical Properties MERCURY Occurrence Physical Prop- 
ertiesChemical Properties MERCUROUS COMPOUNDS Mercurous 
Oxide Mercurous Chloride Mercurous Bromide Mercurous 
Iodide Mercurous Nitrate Mercurous Sulfate Mercurous Sul- 
fide MERCURIC COMPOUNDS Mercuric Oxide Mercuric Chlo- 
ride Mercuric Bromide and Iodide Mercuric Nitrate Mercuric 
Sulfate Mercuric Sulfide Mercuric Cyanide Mercury Fulmi- 
nate Complex Compounds of Mercury Mercuric Ammonia 
Analytical Reactions of Mercury General Relations of the Group. 



General COPPER History Occurrence Metallurgy Electric 
Refining Physical Properties of Copper Chemical Properties 
Alloys Cuprous Compounds Cuprous Oxide Cuprous Chlo- 
ride Cuprous Bromide Cuprous Iodide Cuprous Cyanide 
Cuprous Sulfide The Cupric Compounds Cupric Oxide and 
Hydroxide Cupric Chloride Cupric Sulfate Cupric Nitrate 
Cupric Carbonate Cupric Acetate Cupric Sulfide Copper 
Ferrocyanide, Osmotic Pressure Analytical Properties of Cop- 
per Voltaic Cells SILVER General Occurrence Metallurgy 
Amalgamation Process Leaching Processes Physical Properties 
Chemical Properties of Silver The Oxides of Silver Silver 
Cyanide Silver Nitrate Silver Sulfate Silver Sulfide Silver 
Carbonate Other Salts Photography Mirrors Analytical 
Properties of Silver GOLD Occurrence Metallurgy Amalga- 
mation Processes Smelting Processes Leaching Processes Phys- 
ical Properties The Oxides and Hydroxides The Halogen Com- 
pounds of Gold Sulfides The Complex Cyanides Analytical 
Properties of Gold. 




The Aluminum Sub-group ALUMINUM General Occurrence 
Preparation of the Metal Physical Properties Chemical Prop- 
erties Goldschmidt Process Aluminum Oxide and Hydroxide 
Aluminum Chloride Aluminum Sulfate Mordants Aluminum 
Acetate Aluminum Sulfide Aluminum Silicate Clay Porcelain 
and Pottery Ultramarine Hydraulic Mortars and Cements 



and Metallurgy Physical Properties Chemical Properties 
STANNOUS COMPOUNDS Stannous Oxide Hydroxide and the 
Stannites Stannous Chloride Stannous Sulfide STANNIC COM- 
POUNDS Stannic Oxide and Hydroxide Stannic Chloride Stan- 
nic Sulfide Analytical Properties LEAD Occurrence Metal- 
lurgy Physical and Chemical Properties Oxides and Hy- 
droxides Lead Dioxide Lead' Chloride Lead Nitrate Lead 
Acetate Lead Carbonate and White Lead Lead Chromate 
Lead Sulfate Lead Sulfide The Storage Battery Analytical 
Properties of Lead. 


GROUP V . 426 

SENIC Occurrence Preparation Physical and Chemical Prop- 
erties TRIVALENT COMPOUNDS Arsine Arsenic Trioxide Ar- 
senious Acid The Halogen Compounds Arsenic Trisulfide 
Colloidal Solutions PENTAVALENT COMPOUNDS Arsenic Pen- 
toxide and Arsenic Acid Pentasulfide Divalent Compounds 
Stibine The Trioxide and Its Acids Halogen Compounds 
Other Salts Antimony Trisulfide PENTAVALENT COMPOUNDS 
Antimony Pentasulfide Antimony Pentachloride Antimonic 
Acids and the Pentoxide BISMUTH General Occurrence and 
Preparation Properties Bismuth Trioxide Bismuth Salts 
Bismuth Trisulfide Other Compounds. 



CHROMIUM General Occurrence Preparation Properties 
Chromous Compounds Chromic Compounds Chromic Chloride 


Chromic Sulfate Chromic Acid and the Chromates Bichro- 
mates Chromyl Chloride Photochemical Reactions Perchromic 
Acid Analytical Reactions of Chromium MOLYBDENUM 



Occurrence Properties Divalent or Manganous Compounds 
Manganous Sulfate Manganous Sulfide Manganous Ammon- 
ium Phosphate Manganous Borate Trivalent Manganese 
Tetravalent Manganese Hexavalent Manganese, The Manga- 
nates Heptavalent Manganese, The Permanganates Analytical 
Properties of Manganese. 



THE IRON FAMILY IRON General Occurrence Metallurgy of 
Iron Bessemer Process Open Hearth or Siemens-Martin Process 
Crucible Steel Wrought Iron Physical Properties of Iron 
Chemical Properties of Iron Ferrous Compounds Ferrous Sul- 
fate Ferrous Sulfide Other Ferrous Compounds Ferric Com- 
pounds Ferric Hydroxide Ferric Chloride Ferric Sulfate 
Ferric Thiocyanate Ferric Sulfide Pyrite Ferric Phosphate 
Cyanogen Compounds Ferricyanides Oxalates Ferrates The 
Corrosion of Iron Analytical Properties of Iron COBALT 
Cobaltous Compounds Cobalt Sulfate Cobalt Nitrate Cobalt 
Sulfide Cobalt Glass Cobaltic Compounds Complex Com- 
pounds NICKEL Properties Nickel Compounds Oxides and 
Hydroxides Edison Storage Cell Nickel Carbonyl THE PLATI- 
NUM METALS Ruthenium and Osmium Rhodium and Iridium 
Palladium and Platinum Platinum Platinum Compounds 
Platinous Hydroxide Platinic Hydroxide Analytical Proper- 
ties of Platinum. 


INDEX . 502 



Chemistry is a branch of natural science and deals principally 
with the properties of substances, the changes which they 
undergo, and the natural laws which describe these changes 

It is a very difficult matter to convey thought from one person 
to another by means of words, and anything like accuracy can 
only be attained when the words have as nearly as possible the 
same meaning to each. For this reason it is necessary to discuss 
at some length the significance, in connection with chemistry, of 
some of the terms used in the opening statement. 

Properties of Substances. We are able to perceive objects 
around us. Each of these objects is called a body and the signs 
by which it makes its presence known to us are called the prop- 
erties of the body. When we find ourselves surrounded by a 
number of different bodies we instinctively begin to arrange 
them into groups according to certain points of similarity in 
their properties. We may for example form a class of bodies 
called bottles and group together all objects having the general 
shape of bottles. When we come to examine the different 
members of the class, we find that they possess marked dif- 
ferences in properties, other than those of shape or size. Accord- 
ingly, we at once set up a number of sub-classes such as glass 
bottles, stoneware bottles, rubber bottles, etc., and say that these 
differ because they are made of different substances. Substances 
then, are the things of which bodies are made. As examples of 
substances we may give iron, salt, sugar, lead, etc. 

It now remains to distinguish between the properties of bodies 
and those of substances. This distinction may be brought out 
by considering what we do in forming the sub-class glass bottles. 



In grouping the glass bottles together, we pay slight attention 
to the properties of shape and size, but look to other character- 
istics which we say are the properties of the substance glass. 
Every part of a glass bottle presents the properties of the sub- 
stance glass, and to precisely the same degree as every other part ; 
further, if the bottle be broken into pieces, that body will cease to 
exist and in its place there will be a collection of bodies called 
pieces of glass. If we neglect shape and size, the properties of 
all these bodies are alike and identical with those of the glass 
bottle from which they are formed. The properties of bodies, 
then, aside from those of shape and size are the properties of the 
substance from which the bodies are made. The properties of 
the body, as such, are those of shape and size. A given body 
cannot be a barrel unless it has a certain shape and is of a definite 
size; if it is larger it is a hogshead; if smaller, a keg. 

Changes. Experience shows us that the properties of sub- 
stances can be altered in various ways. 

Every alteration constitutes a change, and for purposes of 
convenience, changes are arbitrarily divided into two classes, 
physical and chemical. 

A physical change is one which alters only a very few of the 
properties of a substance. The moving about of a body or the 
heating of a piece of iron are examples of physical changes. 

A chemical change is one which alters all or nearly all of the 
properties of a substance. In fact, after a substance has under- 
gone a chemical change, we are unable to recognize the presence 
of the original substance, and in its place we find one or more new 
substances with different properties. The burning of wood or 
the rusting of iron are familiar examples of chemical changes. 

From these definitions it would seem to be an easy matter to 
decide whether a given change is physical or chemical. And so 
it is in most cases. However, there are all gradations in the 
number of properties altered in the change, and in some cases 
it becomes really impossible to decide definitely to which class 
they belong. This is due to the fact that there is no real dif- 
ference between the two changes; we have simply arbitrarily 
drawn a distinction as a matter of convenience in our general 
scheme of classification. The freezing of water is an example 
of a change which it is difficult to classify. So many of the 


properties of ice are different from those of water that one is 
inclined to call freezing a chemical change, but on the other hand 
the transformation takes place so easily, simply upon changing 
the temperature, that probably the majority of people consider 
it a physical change. 

In most cases there is no difficulty. For example, a piece of 
rubber when rubbed with woolen acquires the property of attract- 
ing bits of paper, but is not otherwise altered. Without question 
this is a physical change. If the rubber be brought in contact 
with a flame, it will take fire and burn with a smoky flame and 
continue to do so until all the rubber has disappeared. Soon 
after the rubber begins to burn a very strong odor will be noticed 
which must be due to something produced from the rubber. 
Since we have in this case the complete disappearance of the 
properties of rubber and the appearance of the properties of a 
new substance, this is a chemical change. 

Natural Laws. Having discussed the properties of substances 
and the changes which they undergo, we may now turn to the 
natural laws which describe these changes. 

Experience tells us that under like conditions events repeat 
themselves in a very large measure, and that the more nearly 
the conditions are reproduced the more closely are the events 
duplicated. So generally is this the result of our experience that 
we finally become convinced that if it were possible to reproduce 
the conditions exactly, the events would be exactly duplicated 
As a result then of our experience we are able to say, 
after repeated trials, just what takes place under certain con- 
ditions. Such a statement is a law of nature. A law of nature 
differs then essentially from a law of man in that it is simply a 
statement of what does happen and has in it no element of com- 
pulsion. Man is so insignificant a part of nature that he can- 
not presume to dictate to her but can only observe, and learn 
to make the condition such that the operations of nature shall be 
as favorable to himself as possible. A law of nature is a state- 
ment of the way nature works and should be so worded that it 
describes as large a number as possible of single phenomena and 
gives us the maximum amount of information concerning each. 

As an illustration of a law of nature, we may select the law 
of falling bodies. We know as a matter of common experience 


that heavy bodies if unsupported fall to the earth, and a state- 
ment to this effect is a law of nature. It would be much more 
useful if it gave us the results of our experience as to the velocity 
of the bodies after falling for given times and in addition the space 
passed over during certain times of fall. The law might then be 
worded as follows: " All heavy bodies fall toward the earth with 
a velocity which is equal to the force of gravity times the time 
that the body has been falling and the space passed over is equal 
to one-half the force of gravity times the square of the time of 

Obviously from what has been said the wider our knowledge 
of the laws of nature the better equipped we are for life. They 
are all directly or indirectly the results of experience and their 
formulation is one of the most important works of science. 

The Fundamental Law of Chemistry. We find upon examining 
the properties of different bodies composed of the same substance 
that they agree exactly in all their essential properties, that is, 
the properties other than those of shape and size. We find too 
that other substances have radically different properties and 
that the change from one to another is sudden, leaving gaps 
which are not filled in by gradual alterations in the properties. 
We have then as many absolutely distinct sets of unvarying 
properties as there are substances. 

Bodies may be arranged in classes such that the different members 
of each class agree exactly with each other in all their essential 
properties. The different members of each class are the bodies 
composed of the same substance. The law just given is known 
as the law of the definiteness of properties and is often called 
the fundamental law of chemistry. 

Properties of a Substance and the Substance. We have de- 
fined properties as the signs through which objects manifest them- 
selves to us and have spoken of them as though they belonged 
to substances, and. of the substances as though they in some 
way possessed the properties. Indeed the original meanings of 
the words would convey these ideas, and it is very hard to get 
away from them. However, when we come to consider just 
what there is about a substance which is not a property of that 
substance and which might be that which possesses the proper- 
ties we are completely at a loss. Everything that we know about 


a substance is a property of that substance and if by experimen- 
tation we find out anything more, that will also be a property; 
farther than this it seems to be impossible for us to conceive of 
anything concerning a substance that is not a property of that 
substance. One can be readily convinced of this by trying to 
think of anything about a familiar substance that is not a prop- 
erty of that substance. A realization of this fact does not 
make our conceptions of the substance any the less definite 
because these, properties are the real things about the substance, 
which we can know and measure. In fact our idea of the sub- 
stance is simply the sum of all these properties which we know. 
For us then a substance is simply a specific group of essential 
properties which always occur together*and to an unvarying degree 
under given conditions. This may be used as a working definition 
of a substance. If there is anything more to a substance than 
its properties we can know nothing of it. We cannot even 
imagine anything about it. So we will leave the question of the 
actual existence of a possessor of properties to speculative 
philosophy, and in matter of fact chemistry when we say 
substance we will mean properties. 

Every known substance has a name which in a way stands 
for the properties of the substance. These names can mean to 
us only as much as we know of the properties of the substances. 
So in studying chemistry we must take care that we do not merely 
learn the names of the various substances with which we deal, 
but also that we make these names mean something to us by 
learning the more important properties of the substances. 

Physical and Chemical Properties. It is convenient to divide 
properties into two classes physical and chemical. A physical 
property Is one which can be detected and measured without 
causing the substance to undergo more than a physical change. 
As examples we might give color, density, conductivity for 
heat or electricity, etc. 

A chemical property is one which is only revealed when the 
substance is transformed into something else and consequently 
undergoes a chemical change. One of the properties of sulfur is 
that it burns with a pale blue flame, and finally all disappears, 
leaving behind something which is invisible but which has a very 
strong smell. Since this property is shown only when the sulfur 


is transformed into a new substance, this is, therefore, a chemical 

Since there is no real difference between physical and chem- 
ical changes, there is none between the two sets of proper- 
ties, but nevertheless, it is convenient to make the arbitrary 

Identification of Substances. The chemist is very often con- 
fronted with the problem of deciding as to whether two different 
bodies are composed of the same or of different substances. The 
decision rests upon the answer to the question, Do the bodies 
have exactly the same essential properties? If they do, they are 
composed of the same substance; if they do not, of different 
substances. Evidently the question can only be answered after 
carefully investigating the properties. To be perfectly sure, it 
would seem to be necessary to compare all of the properties 
because two substances might agree in most of their properties 
and yet differ enough in some to make them different substances. 
The labor required for the comparison of all of the essential 
properties is so great that it is never done. The chemist com- 
pares some of the essential properties, and if these agree exactly 
he decides that the two bodies are composed of the same sub- 
stance. In doing this he takes advantage of a law which states 
that if two bodies agree exactly in some few of their essential 
properties they will agree exactly in all and are composed of the 
same substance. 

The Characterization of a Substance. The. properties chosen 
for investigation in order to characterize a substance vary with 
the case, but naturally they are, in general, those which can be 
most readily observed and measured, or else they are the ones in 
which the substance differs most from other substances. 

The impressions produced upon our sense of sight, taste, and 
smell can be very easily determined, and are but rarely omitted. 
The physical state (solid, liquid, or gaseous) of the substance at 
ordinary temperature and pressure and the conditions under 
which it changes from one state to another are easily determined 
and important properties. The solubility of a substance in 
water is another valuable characteristic. A few words concern- 
ing some of these frequently studied properties would seem to be 
in order here. 


Physical States. We distinguish three different ways in which 
substances fill space and call these the physical states, giving 
them the names solid, liquid, and gaseous states. 

The Solid State. A substance is a solid if a given body com- 
posed of this substance has both a definite volume and a definite 
shape of its own independent of its surroundings. 

The Liquid State. A substance is a liquid if a given body com- 
posed of this substance has a definite volume but takes its shape 
from its surroundings. A liquid always runs down to the lower 
part of the vessel in which it is placed and takes on the shape of 
the vessel in so far as it can fill the container. The free surface 
of a liquid tends to be flat and parallel to the surface of the earth 
unless the body of the liquid is small, when the surface is rounded. 
If the body of liquid be very small and the surface entirely free, 
the liquid takes on the form of a sphere modified more or less by 
the action of gravity. 

The Gaseous State. The gaseous state is characterized by the 
fact that a gaseous substance has neither definite volume nor 
shape, but always fills the container in which it is placed and 
consequently takes its shape and volume from the surroundings. 

It is a very easy matter to tell, in the great majority of cases, 
whether a given substance is solid or gaseous. There are, how- 
ever, a few substances which are like wax in that they keep a 
fairly definite shape for some time and so give the impression that 
they are solids. If left for a longer time they slowly change their 
shape and flow, thus showing that they are really viscous liquids. 
But such substances are comparatively rare and give but little 
trouble in classification. 

The physical state of a substance may vary with the conditions 
such as temperature, etc. With rising temperature, solids always 
tend to become either liquid or gaseous, and liquids to become 
gaseous. With falling temperatures gaseous substances always 
tend to become either liquids or solids, and liquids to become 

Solubility and Solution. Many substances when brought in 
contact with water mix with the latter in varying proportions 
and form a homogeneous liquid which has in a general way the 
properties of the water and of the other substance. Such a 
homogeneous mixture, i.e., one in which every distinguishable 


particle is exactly like every other particle in essential proper- 
ties, is called a solution. The ability of a substance to form a 
solution is called its solubility and is an important property of 
the substance. A more extended discussion of solution will be 
given at a later point, where it will be shown among other 
things that solutions are not necessarily liquid but may be solid 
or gaseous as well. 

Mechanical Mixtures. Substances are often present together 
in such a way that they do not form a homogeneous mixture but 
the different parts may be distinguished and more or less easily 
separated. Such mixtures are called mechanical mixtures. 
Muddy water may be given as an example of such a mixture. 
Examine it with a lens and the particles of silt can be seen. 
Allow it to stand or pass it through a filter and the silt may be 
separated from the water. Muddy water is not homogeneous 
and hence is not a solution. Solutions stand between mechanical 
mixtures and a very important class of substances, with which 
we will become familiar in a little time, known as chemical 


Units. Some of the properties of substances, as for example 
color, can be ascertained by simple inspection, while others re- 
quire more or less elaborate experiments to bring them to light. 
The constant aim is to represent these properties by numbers. 
To do this a unit must be decided upon and the property carefully 
measured in terms of this unit. Almost all of the units used in 
scientific work are derived directly or indirectly from three 
fundamental units, the centimeter, the gram, and the second, 
and the whole scheme of units is called the C. G. S. system. The 
second is the unit of time. It is in use in daily life and is familiar 
to everyone. 

Length. The centimeter is the unit of length and is the 
one-hundredth part of the length of a certain bar of platinum 
carefully preserved in Paris, which is called the "Standard 
Meter." This standard meter was intended to be the 
Ko, ooo, ooo of the earth's quadrant or quarter circumference, 
measured on the meridian of Paris, but afterward turned out 
to be something different from this owing to an error in the 
measurement. An inch is equal to a little more than 2.5 cm. 

Mass. The gram is the unit of mass. It is the K.ooo P ar ^ 
of the mass of a certain piece of platinum called the standard 
kilogram. It was intended that the gram should be the mass 
of a cubic centimeter of water at its temperature of maximum 
density, 4C., but actually the mass of the latter is 0.999982 
grm. This differs so slightly from unity that the mass of 1 c.c. 
of water at 4C. is usually taken as exactly 1 grm. 

Weight. A heavy body is always pulled toward the earth 
with a certain force. This force is called its weight and is 
proportional to the mass of the body. The weight of a given 
body varies with its position on the earth, but an average is 981 
dynes for each gram of the body's mass. A dyne is that force 



which acting on 1 gram for 1 second will generate a velocity of 
1 cm. per second. 

If one takes as the unit of weight, not the dyne, but, as is 
commonly done the weight of 1 grm., then the numerical value 
of the weight of a body in terms of this unit is the same as 
the mass of that body. The statement that a body has a weight 
of 10 grm. means that its weight is the same as that of 10 grm., 
9,810 dynes say, and consequently the mass is 10 grm. 

Volume. The unit of volume is the cubic centimeter or the 
still larger unit, the liter, which is 1,000 c.c. 

Energy. Work is usually denned as the product of a force and 
the distance through which it acts : energy as work or the ability 
to do work. It has been further shown that there are many forms 
of energy such as heat, light, electrical energy, chemical energy, 
etc., and that these forms may be changed one into another 
without loss. In fact in an isolated system of bodies, energy 
can be neither created nor destroyed, but remains present in un- 
changed amount. This is the law of the conservation of energy. 

The unit of energy is the erg, and this is the work done by a 
force of 1 dyne acting through a distance of 1 cm. This unit 
is rather small, and so we have the larger units, the joule, which 
is 10,000,000 ergs, and the kilojoule which is 1,000 joules. 

Whenever work is done upon a body, the latter acquires the 
ability to do work, and hence an increase in energy. The 
simplest way to regard this change of work into energy is to 
consider work as one of the forms of energy. If this is done 
energy may be defined as work and everything else which may be 
obtained from work and reconverted into work. 

Conservation of Weight and of Mass. One of the most useful 
instruments to the chemist is the balance. This is a device by 
means of which the weights of bodies may be compared, and 
since the weights are proportional to the mass, the relative 
masses may be determined. 

If we ask the question, how are the masses of substances 
affected when they undergo chemical change, we can obtain the 
answer by weighing the reacting substances before and after the 
change. We will find that in many cases there are apparent 
losses in weight, as in the case of burning a piece of wood and in 
other cases a gain in weight, as in the rusting of iron. But these 


changes take place in contact with air and hence the changes in 
weight may be due to something escaping to the air in the case of 
the wood, or coming in from the air in the case of the iron. 

This possible source of change in weight may be excluded by 
sealing the whole reacting substance air-tight in a flask. When 
this is done, it will be found that whatever changes in weight may 
be observed will be inside the limits of experimental error, for 
all but the most refined experiments, and even in these cases the 
small changes observed may perhaps be due to some source of 
error as yet unknown. 

Neither is there any known physical change which will altel 
the weight at any given spot of the substance undergoing the 
transformation and therefore we at once set up the law of the 
conservation of weight which is, that neither a physical nor a 
chemical change alters the total weight of the system undergoing 

Since mass is proportional to weight we may get from the above 
the law of the conservation of mass. The total mass of a system 
is the same before and after any chemical or physical change which 
may take place within the system. This simply means that 
although chemical changes alter so radically the other properties 
of substances they do not change that property called mass. 

Density. The density of a substance may be defined as the 
ratio of the mass of a given body composed of that substance 
to the volume of the body. It is expressed in terms of grams 
per cubic centimeter. 

Specific Gravity. The specific gravity of a substance is the 
ratio of the weights of equal volumes of that substance and of 
some other substance taken as a standard. Water is usually 
taken as the standard. The temperature of both substances 
must be stated^ in order that specific gravity may be definite. 
If water at 4C. is taken as the standard substance, since at that 
temperature its density is one, the numerical value of the specific 
gravity will be the same as that of the density, although the 
conceptions are entirely different. 


Discovery of Oxygen. Oxygen was discovered August 1, 1774, 
by the English chemist Priestly who prepared it by heating 
mercuric oxide. It was independently discovered something like 
a year earlier than by Priestly, by the Swedish chemist Scheele, 
but his results were not published until 1777. Scheele obtained 


- -Oxygen 

FIG. 1. 

it by heating saltpeter, mercuric oxide, manganese dioxide, and 
a number of other substances. 

Priestly's original experiment, in a somewhat modified form, 
may be easily repeated. Mercuric oxide, which is a reddish 
powdery substance, is placed in a doubly bent tube (Fig. 1) which 



must be of hard glass. It is then strongly heated, and soon the 
mercuric oxide darkens in color and becomes almost black. If 
the heating be continued a bright metallic film appears on the 
cooler part of the tube, and at the same time, if the end of the 
delivery tube be placed under water, bubbles of gas will be seen 
to rise through the water. This gas may be collected by filling 
a vessel with water and placing it mouth downward in a trough 
of water. Upon bringing the mouth of the vessel over the 
delivery tube, the bubbles of gas will pass up into the vessel and 
gradually displace the water. When the vessel is full of the gas 
it may be removed from the water and placed mouth upward 
on the table for experimentation, taking care to keep it covered 
with a glass plate. It will be as transparent as air and in fact 
will appear just like air. It may, however, be distinguished 
from air by its action toward a feebly glowing spark on the end 
of a splinter. In oxygen the spark bursts into flame, while in 
air it barely continues to glow. 

If the process of heating the mercuric oxide be continued, the 
oxide will be seen to diminish gradually in quantity and the 
bright metallic film to increase, and all the while the gas oxygen 
will be given off. This will go on until 'the mercuric oxide has 
all disappeared, then the process comes to an end. If the bright 
metallic film be examined it will be found to consist of drops of 
mercury, which is often called quicksilver. So we have the one 
substance, mercuric oxide, decomposing into two substances, 
oxygen and mercury which have entirely different properties 
from the mercuric oxide; therefore, this is a chemical change. 

Combustion. Experience from infancy has made us familiar 
with that most important set of chemical phenomena called com- 
bustion. We see the coal or wood burn in the stove, the oil or gas 
in the lamps. We know that a great amount of heat is universally 
given off during such a process. In most of the above-mentioned 
cases, it is very apparent that the burning substance decreases 
in weight. There are, however, other cases which we would 
certainly classify as combustion in which the weight is increased 
during the process of burning. For example, if a piece of metallic 
magnesium, in the form of a ribbon, be held in the flame of a 
match it will take fire and burn with an exceedingly bright light. 
The metal disappears and in its place is left a white powdery sub- 


stance which weighs two-thirds more than the metal did before 
burning. Finely divided iron will burn when brought into a 
flame, forming a black brittle substance which weighs more than 
the original iron. In these and other instances of the same kind 
which might be mentioned, the products of combustion are solids, 
and hence easily obtained in a weighable form. This suggests 
that perhaps in every case of combustion the products are really 
heavier than the combustible substance, and that in the case of 
the wood, etc., first mentioned, the products are gaseous and 
hence escape our observation. That this is so may be easily 
shown in the case of a lamp by holding a cold object over the 
burning lamp when it at once becomes covered with a film of 
liquid water. Since this water will not accumulate if the lamp 
is not lighted, we conclude that the water is one of the products 
of the combustion of the oil. If a drop of lime water, on the end 
of a glass rod or a loop of wire, be brought over the flame of a 
lamp it will quickly become clouded. Now water vapor will not 
produce this cloud in the lime water nor will it appear unless 
the lamp is burning, and therefore we conclude that something 
besides water is formed during the burning of the oil. One of 
the characteristic properties of carbon dioxide is its ability to 
cloud lime water, and we then conclude that carbon dioxide as 
well as water is formed during the combustion of the oil. Proc- 
esses by which the presence of certain substances may be de- 
tected are called tests or reactions, and the substances used, 
such as the lime water in this case, are called reagents. By 
applying these. same tests to burning coal, wood or gas, it will 
be found that in these cases also, carbon dioxide and water vapor 
are formed. If the coal is anthracite, comparatively little water 
vapor and much carbon dioxide will be produceol. 

There is a solid white substance called caustic potash which 
has the power to take up both water vapor and carbon dioxide, 
and to hold them in such a form that they may be weighed. 

If a glass cylinder loosely filled with this substance be placed 
over an unlighted candle and the whole counterpoised upon a 
balance and the candle then lighted (Fig. 2), the ca\istic potash 
will absorb the water vapor and carbon dioxide produced during 
the burning, and it will be very quickly seen that the weight of 
the whole is increasing, although the candle is visibly becoming 



smaller. In some such a way as this it may be shown that the 
weight of the products of combustion is always greater than that 
of the combustible substance. Since in every case the products 
of combustion have different properties from the original sub- 
stance, the changes that take place are chemical changes. 

FIG. 2. 

The law of the conservation of weight states that the weight 
of the products of a chemical change is the same as that of the 
substances before the change. We must conclude then that there 
is some other substance than that which burned taking part in 
the chemical change which we call combustion, and that the 
weight of this other substance is equal to the difference in weight 
between that of the products of combustion and the combustible 

Role of the Air. All these changes which we have been con- 
sidering take place in contact with the air, so it is very probable 
that this other substance is present in the air. This is con- 
firmed by the fact that it is impossible to make a combustible 
substance burn in a vacuum, and further, when it burns in a 
limited amount of air a part of the air is used up, or disappears, 
but never more than one-fifth of the total volume of the air. 
After one kind of combustion has used up one-fifth of the air, no 



FIG. 3. 

other substance will burn in the remaining four-fifths. This 
indicates clearly that the air consists of at least two gaseous 
substances, one present to the extent of one-fifth of the total 
volume; and taking a most important part in the phenomenon 
of combustion. 

If mercury be heated for a number of days to a temperature 
a little below its boiling-point in contact with a limited volume 
of air, a reddish substance, mercuric oxide, is formed on the 

mercury, and the air gradually 
decreases in volume. If the 
mercuric oxide obtained in this 
way be heated as described on 
p. 12, it yields mercury and 
oxygen, and the volume of oxy- 
gen so produced is just equal to 
the decrease in the volume of 
the air. This oxygen has the 
power of supporting combus- 
tion to an extraordinary degree 
and in the burning, all the oxy- 
gen may be used up. The residual air will not support combus- 
tion, but if the oxygen obtained by decomposing the mercuric 
oxide be added to it, the air regains all its original properties. 

These experiments were first performed by the French chemist 
Lavoisier shortly after the discovery of oxygen, and had much 
to do with our present conception of combustion. The apparatus 
which he used is shown in Fig. 3. 

Combination and Decomposition. We have seen above that 
mercuric oxide may be transformed into mercury and oxygen. 
This process is an example of what is called decomposition, and the 
reversed process, the formation of the mercuric oxide from oxygen 
and mercury, is an example of combination. Processes like these 
compose a large part of the changes studied in chemistry. 

Preparation of Oxygen. Although free oxygen is present in 
the air in such enormous quantities, the problem of separating it 
from the substances with which it is mixed is by no means simple. 
(See liquid air, pp. 25 and 200.) 

There are many other substances besides those already given 
which will yield oxygen upon being heated. Among these may 



be mentioned potassium chlorate and potassium permanganate, 
and in addition oxygen may be obtained by the decomposition 
of water by the electric current and by acting upon sodium per- 
oxide with water or dilute acid. The method most in use in the 
laboratory is that of heating potassium chlorate. This is a white 
crystalline substance which melts at about 360C. and decom- 
poses at a higher temperature into oxygen and potassium 
chloride. The rate of decomposition increases rapidly as the 
temperature is raised. Curiously enough if it be mixed with 

FIG. 4. 

something like one-third its weight of manganese dioxide the 
decomposition takes place much more rapidly than with the 
potassium chlorate alone, being quite active at as low a tempera- 
ture as 200C., and after the operation the manganese dioxide 
is found to be unaltered except, perhaps, that the particles may 
be a little finer than before. Although manganese dioxide will 
yield oxygen at a high temperature, it will not do so in noticeable 
quantity at 200C., and its action here seems to be to greatly 
increase, by its mere presence, the rate at which the potassium 
chlorate decomposes. Actions of this kind are common and are 
known as catalytic actions. They will be considered more in 
detail later. 


The decomposition of the potassium chlorate when mixed 
with manganese dioxide may be carried out in a hard glass flask 
(Fig. 4). The oxygen is collected by the displacement of water 
as explained above. 

Physical Properties. The gas so obtained is colorless, odorless, 
and not very soluble in water or it could not be collected over 
this liquid. 

One of the important physical properties of every substance 
is its density. The simplest way to determine the density of 
oxygen is to decompose a known weight of either mercuric oxide 
or potassium chlorate and weigh the residue. The loss in weight 
is the weight of the oxygen. If the unit of weight is the weight 
of 1 grm. the loss in weight gives at once the mass of the oxygen ; 
and to obtain its density it will only be necessary to measure the 
volume of the oxygen, and divide the mass by the volume. For 
example, 1.6000 grm. of mercuric oxide gave 0.1111 grm. of oxygen 
which measured 87.15 c.c. under a pressure of 74 cm. of mercury 
and at 25C. The density of oxygen under these conditions is 
therefore - 11]L M7.i5 = 0.0012752 grm. per c.c. 

Standard Conditions. If the volume of a gram of oxygen, or 
of any other gas for that matter, be determined under different 
atmospheric conditions it will be found to vary widely. Other 
things being the same, the higher the temperature, the greater 
the volume, and the higher the pressure the smaller the volume. 
It can be readily seen from this that the density of a gas is not 
definite unless the volume of a given mass of the gas be measured 
at a certain definite temperature and pressure. For the sake of 
uniformity it is the universal custom to take the temperature of 
melting ice, or 0C., as the standard temperature since this can be 
so easily reproduced. As the standard pressure, is taken the 
mean pressure of the air at the sea level, which is equal to that of 
a column of mercury 76 cm. in height. Since the density of the 
mercury changes with the temperature, the mercury should also 
be at 0C. 


Boyle's. As it is usually not convenient to measure the vol- 
ume of gases under standard conditions, it is a matter of great 
importance to know just how they change their volumes with 


changes in pressure and temperature, so that their volumes may 
be read under any convenient conditions and corrected to what 
they would be under standard conditions. Fortunately it has 
been found that gases behave so nearly alike that the same cor- 
rections will do for all unless extreme accuracy is demanded. 

The simplest relationship is that between the volume and the 
pressure, and this has been known ever since 1660, when it was 
discovered by the English scientist Robert Boyle. 

Boyle found that if the pressure upon a fixed mass of a given 
gas, at constant temperature, be doubled, trebled, or quadrupled, 
the volume of the gas would be reduced to one-half, one-third, or 
one-fourth of 'its original volume, or in other words the volume of 
a fixed mass of a given gas at constant temperature is inversely 
proportional to the pressure. This is known as Boyle's law, and 
may be represented as follows: 

v = k- (1) 


when v and p are the simultaneous values of the volume and 
pressure under which the gas is measured. 

The above relation may also be expressed in the following 

= p 2 v 2 = pv = const. (2) 

where piVi and p 2 v 2 are pairs of simultaneous values of pressure 
and volume. In words this is expressed as follows: The product 
of the simultaneous values of the pressure and the volume for a 
fixed mass of a gas at constant temperature is a constant. This is 
simply another way of wording Boyle's law. 

As an example of the practical application of this law we may 
take the data given above for the density of oxygen. There we 
had 87.15 c.c. of oxygen at 25C. and a pressure of 74 cm. of 
mercury. What would be the volume at the same temperature 
but under the standard pressure of 76 cm. of mercury? 

Taking equation (2) we have 

p 2 V 2 

and letting p 2 = 74, v 2 = 87.15 c.c. and pi = 76 cm. of mercury, 


vi = volume at this standard pressure. Solving for vi we get: 


74 X 87.15 
vi = - ^r- - = 84.85 c.c. 

This would be the volume of the gas if measured at the tempera- 
ture 25C. under the standard pressure. 

In order to calculate the volume which the gas would occupy 
at the standard temperature as well as pressure, it would be 
necessary to ascertain the influence of a given change in the tem- 
perature upon the volume of a fixed mass of a gas at constant 
pressure, and apply the proper correction to the calculated 
volume of the gas at the standard pressure. As indicated above, 
the volume of a fixed mass of gas increases with rising tempera- 
ture, the pressure remaining constant. If we take the volume 
occupied by a certain mass of a gas at the melting-point of ice or 
0C. as unity, and then heat the gas to the temperature of water 
boiling under a pressure of 76 cm. of mercury, the volume of the 
gas will be found to be 1.367, so that the increase has been 0.367. 
If now we divide the difference in temperature between the 
melting-point of ice and the boiling-point of water into 100 equal 
parts, as has been done in the Centigrade thermometer system, 
the increase in volume for each degree is 0.00367 or J^?s of the 
volume at 0C. This is called the coefficient of expansion of the 
gas and is represented by alpha, "a." 

The law describing this behavior is called Gay Lussac's and 
Dalton's Law or sometimes Charles' Law and was discovered 
simultaneously by Gay Lussac and Dalton in 1802. It may be 
expressed as follows : The volume of a fixed mass of a gas at con- 
stant pressure is increased by ^73 of its volume at Centigrade 
for each degree's rise in temperature.. 

The mathematical expression is 

v t = (1 + a t)v 

where v = volume at 0C., v* = the volume at the temperature 
t, t = the temperature, and a the coefficient of expansion, 
0.00367 or ^73. The conditions under which this will hold 
are that both the mass of the gas and the pressure upon it 
must remain constant. 




Absolute Temperature. Since all gases are so regular in their 
behavior toward changes of temperature the change in the 
volume of a fixed mass of a gas under constant pressure has 
been taken as the measure of the change in temperature. Gas 
thermometers are really the standard instruments for the 
measurement of temperature at the present time. Volume Temp . 
Mercury thermometers are used as working instru- 
ments because they are somewhat more convenient 
than gas thermometers. 373-10 iooc. 

In the Centigrade thermometer system, the 350 
melting-point of ice has been called and the 
boiling-point of water 100. If a fixed mass of a 
gas at constant pressure be raised in temperature 300 
from the melting-point of ice to the boiling-point 
of water (Fig. 5), its volume will be increased by 
10 %73 of its volume at 0C. Since we have 25 
already decided to call the temperature interval 
between the freezing- and the boiling-point of 
water 100 and to take the change in the volume 
of a gas as the measure of the change in the tem- 
perature, we say that if a fixed mass of a gas at lgo 
constant pressure changes its volume by ^73 of 
its volume at 0C. its temperature has been changed 
by 1C. If it changes its volume by 2 %?s of its 100 
volume at 0C. the change has been 20. 

It is a very simple matter to cool a gas to a tem- 
perature below 0C. The gas contracts during the 15 
process and continues to do so as it is cooled to 
lower and lower temperatures. In accordance with 
what has been said above, we consider that the gas FIQ ~ 5 ? 
has been jcooled 1 for each ^ 73 of its volume at 
0C. which it has lost, the pressure and the mass of gas being 
constant. It is very obvious that a gas cannot lose more than 
27 M73 of ^s volume because this would reduce it to zero 
volume. Therefore, the lowest temperature to which a gas 
can be cooled is 273 Centigrade degrees below the Centigrade 
zero and this temperature is called absolute zero. As a matter 
of fact all gases are liquefied or solidified before they reach 
this temperature, and since the gas laws do not describe the 


behavior of liquids or solids the substances would doubtless have 
some volume even if cooled to absolute zero. The lowest tem- 
perature yet reached is about 3 above absolute zero which 
is the temperature at which liquid helium will boil under 
diminished pressure. 

Temperatures reckoned upward from this absolute zero are 
called absolute temperatures. The absolute temperature of 
melting ice is 273 A., and of boiling water 373 A. Tempera- 
tures on the Centigrade scale may be readily changed to the 
absolute scale by adding 273 to the Centigrade temperature. 
Absolute temperatures are usually represented by T and Centi- 
grade by t, so, 

T = t + 273 

There are two important advantages of the absolute over the 
Centigrade system. First, by its use one does away with all 
negative temperatures and second, we are enabled to state the 
law of Gay Lussac and Dalton in the following exceedingly simple 
manner. The volume of a fixed mass of a gas at constant pressure 
is directly proportional to the absolute temperature. In symbols 
this law is expressed as follows: 

v = k'T (3) 

where v is the volume, k' the proportionality factor, and T the 
absolute temperature. 

By a simple transformation of equation (3) we get 

or the volume of a fixed mass of a given gas divided by the 
absolute temperature at which it is measured is equal to a 
constant. From this it follows that 

* = ^ (4) 

To T! 

in which V , T and Vi, TI, are any different sets of simultaneous 
values of the volume and temperature for the same sample of 

With the aid of equation (4), we may readily calculate the 
volume of a given sample of gas at one temperature from its 


volume at another provided the pressure remains constant. This 
may be illustrated by calculating the volume of the oxygen 
obtained on p. 18 at the standard temperature of 0C. or 273A. 
On p. 20 we made the correction for pressure according to 
Boyles' law and found that at 25C. or 298 A. and under the 
standard pressure, the volume would be 84.85 c.c. Letting V 
be the volume at 273A., Vi be 84.85 and TI be 298A., we get 

_Vo_ _ 84.85 
273 298 

V = 84.85 X = 77.74 c.c. 

To obtain the density of oxygen under standard conditions all 
that remains to be done is to divide the mass of the oxygen, 
0.1111 grm. by this calculated volume, 77.74 c.c. 

Density = ^^ = 0.0014293 grm. per c.c. 

The Combined Gas Laws. In calculating the volume of the 
oxygen obtained on p. 18 we first corrected for pressure by 
multiplying the measured volume 87.15 c.c. by the ratio of the 
measured pressure 74 cm. to the standard pressure 76 cm. and 
then took this result, 84.85 c.c., and corrected for temperature 
by multiplying it by the ratio of the standard temperature, 273A. 
to that at which the gas was measured 25C. or 298A. Ob- 
viously these operations might just as well have been combined 
and the double correction made at once, as follows: 
74 27*} 

V = 87.15 X 41 X ~ = 77.74 c.c. at 0C. and 76 cm. 

of mercury. 

This is a perfectly general relationship and may be represented 


V. = Vi X ~ X ^ (5) 

f ll 

We may then calculate the volume of any given sample of 
gas at any desired condition of temperature and of pressure 
from its volume at any other temperature and pressure by 
multiplying the given volume first by the ratio of the pressures 
and second by the ratio of the absolute temperatures. In 


applying this relation it is, of course, necessary to get the proper 
temperature and pressure in the numerator. To do this, it is 
merely necessary for the calculator to ask himself, "Will the 
given change in pressure increase the volume?" If so, the 
larger pressure is put in the numerator; if not, in the denom- 
inator; and correspondingly for the temperature. The use of 
absolute temperatures is advised but may be avoided in the cor- 
rection to standard conditions by using the following equation 
which is easily derived from equation (5) 


V = 

76(l + 0.00367t) 

in which t is the Centigrade temperature at which the volume Vi 
is measured. 

Other Physical Properties of Oxygen. Solubility is one of 
the very important properties of a substance, a knowledge of 
which goes to make up our conception of the substance. As has 
been pointed out above, oxygen is not very soluble in water or it 
could not have been collected over this liquid. Its solubility in 
water, however, is easily measured since one volume of water 
at 0C. dissolves 0.04890 volumes of oxygen. At higher tempera- 
tures the solubility is smaller. Many metals absorb noticeable 
quantities of oxygen. Melted silver, for example, takes up 
something like 10 times its own volume of oxygen, the greater 
part of which it gives off on cooling. The escaping gas causes 
protuberances of silver to grow out from the bead, and produces 
the phenomenon known to assayers as sprouting. 

Liquid Oxygen. If oxygen under atmospheric pressure be 
cooled below 182.5C. it becomes a pale blue liquid which 
has a density of 1.12 and boils under atmospheric pressure at 

182.5C. If the pressure be lowered the boiling-point falls 
until under a pessure of 0.75 cm. of mercury the liquid boils at 

211.2C. Upon raising the pressure the boiling-point rises, 
and finally under a pressure of 50.8 atmospheres the boiling- 
point is 118.5C. If the temperature be raised any higher than 
this the oxygen refuses to remain in or take on the liquid state 
no matter how great the pressure. Hence to liquefy oxygen it 
must be cooled to at least 118.5C., then upon the application 
of sufficient pressure, which at 118,5C, is 50.8 atmospheres 


but is lower if the temperature is lower, the liquid appears. At 
ordinary temperatures or at any temperature higher than 

118.5C. it is not possible to liquefy oxygen by any pressure 
however great. 

All other gases behave as oxygen does in this respect. Each 
gas has a certain temperature above which it is impossible to 
liquefy it. This temperature is called the critical temperature, 
and the pressure which will be just great enough to bring about 
the liquefaction at the critical temperature is called the critical 
pressure. In the case of oxygen, the critical temperature is 

118.5C. and the critical pressure is 50.8 atmospheres. An 
atmosphere is a pressure equal to 76 cm. of mercury. For other 
gases, the critical constants are different, and are characteristic 
for each substance. In learning the properties of any given 
gaseous substance, a fairly definite notion of its critical constants 
and boiling-point should be obtained. 

Air is a mixture of approximately one part by volume of 
oxygen and four parts by volume of other gaseous substances, 
chiefly nitrogen. Liquid air has something like the same com- 
position but is a little richer in oxygen because this is more 
easily liquefied than the nitrogen. If the liquid air be allowed to 
boil away, the nitrogen will come off in relatively greater quan- 
tities than the oxygen and the liquid will gradually become 
richer and richer in oxygen, and finally will yield a gas containing 
95 per cent, or more of oxygen. This is the basis of the most 
important method for the preparation of oxygen upon a com- 
mercial scale. The gas is put upon the market in strong steel 
cylinders under a pressure of 100 atmospheres. 

Both gaseous and liquid oxygen are attracted by a magnet 
although the attracting force is very much smaller than that 
exerted upon an equal weight of iron. 

Solid Oxygen. When oxygen is cooled with liquid hydrogen, 
it freezes to a light blue solid which melts at 227 C. 

Chemical Properties. The ability to support combustion to an 
extraordinary degree is the chief chemical property of oxygen, 
and is made use of in its identification by introducing into 
some of the gas a glowing spark on the end of a stick, 
when it at once bursts into flame. This is the test or reaction 
for .oxygen and the splinter or spark is the reagent. 


There are two good reasons why the wood burns faster in the 
oxygen than in the air. The first is that since air is only one- 
fifth oxygen, more oxygen is in contact with the wood at a given 
time in the pure gas than in air, and naturally they combine 
more rapidly. Second, in the case of the combustion in the air, 
the heat given out in the burning of the wood is divided between 
the wood, the oxygen, the rest of the air and the products of 
combustion; and so of course the temperature cannot rise as 
high as in the case of the combustion in pure oxygen, where it 
has only the wood, the oxygen, and the products of combustion 
to warm. Now the same thing is true in this case that was 
mentioned about the decomposition of potassium chlorate, that 
the reaction goes on much faster the higher the temperature. 
Therefore the rapid combustion in oxygen is due to the more 
abundant supply of the gas and to the higher temperature. 

From this explanation we would expect that other things would 
burn more rapidly in oxygen than in air, and upon trial we 
find our expectation realized. For example sulfur burns in the 
air with a pale blue flame of very little luminosity, while in 
oxygen it burns with a bright blue flame. Phosphorus burns 
in air with a yellowish-white flame, in oxygen the combustion 
is exceedingly rapid and the light is so intense as to be fairly 

Many substances which will not readily burn in the air will do 
so in oxygen, for example iron. If a piece of iron be heated in 
the air, it becomes covered over with a coating of something 
which may be separated rather easily from the iron, and at the 
same time that the coating is formed the iron increases in weight. 
If a thin piece of iron, say a watch spring, be heated white hot in 
oxygen it will take fire and burn vigorously, throwing off sparks 
in a very spectacular manner. The substance left after the com- 
bustion is the same as that formed by heating the iron in the air. 

We know as a matter of practical experience that in order to 
make things burn rapidly it is necessary to raise their tempera- 
ture, and therefore it might seem natural to ask what the kin- 
dling temperature is in each case. Experiment will tell us that 
combustible substances combine with oxygen at temperatures 
far below that at which they burst into flame; that with rising 
temperature the rate at which the combination or combustion 


takes place increases very rapidly until finally the substance 
kindles into flame. It is not possible to locate very accurately 
this temperature at which kindling into flame takes place as it 
depends upon many factors, such as the state of subdivision of the 
substance, etc. Still less is it possible to locate a temperature 
above which a substance will combine with oxygen, while below 
it the reaction will not take place. The rate simply becomes 
smaller and smaller as the temperature falls. We therefore must 
conclude that all combustible substances combine with oxygen 
at all temperatures, although the rate may be immeasurably 

The combination of oxygen with a combustible substance at 
temperatures below the kindling point is called slow oxidation. 
This is not essentially different from combustion, and with rising 
temperature gradually merges into the latter. As much heat is 
given off during slow oxidation as during the rapid combustion 
of the same quantity of the substance. If this heat is retained 
and goes to raise the temperature of the body, as for example, 
when the reaction takes place in the midst of a heap of coal, the 
temperature may rise high enough for the substance to burst 
into flame. When this happens we say it is a case of spontaneous 

Quantitative Relations. If we study with care the decomposi- 
tion of mercuric oxide, taking into account the weights of the 
various substances, we will find that the weight of the oxygen 
formed plus that of the mercury is just equal to the weight of 
the mercuric oxide taken. This shows, of course, that nothing 
else than mercury and oxygen is formed from mercuric oxide. 
We will find too that, no matter how or where the mercuric 
oxide is prepared, there will always be a constant ratio between 
the weight of the mercury and the weight of the oxygen obtained 
by decomposing mercuric oxide. This ratio is also the same as 
that in which the mercury and oxygen combine to form 
mercuric oxide. 

This is simply one example of the law of constant or definite 
proportions, which is one of the most important of the laws of 
chemistry, and may be stated as follows: " Whenever two or more 
substances combine to form another substance, they do so in a 
perfectly fixed and invariable ratio by weight." 


Elements. As we have seen above, mercuric oxide yields two 
substances, oxygen and mercury, whose combined weights are 
equal to that of the mercuric oxide, while their individual 
weights are less. This was spoken of as decomposition, and it 
is a common characteristic of all decomposition processes that 
substances are produced whose individual weights are less than 
the weight of the substance from which they were formed. If 
we examine the chemical transformations which oxygen under- 
goes, we will find that in no case is there a substance formed 
whose weight is less than that of the oxygen used up. In one 
case, that of the formation of ozone (p. 29), the weight of 
the resulting substance is the same as that of the oxygen 
consumed; in all other cases it is greater. We must conclude 
then that oxygen has never been decomposed, that is, changed 
into something having a smaller weight. The same is true for 
mercury, and for about 80 other substances. These substances 
are called the elements, 1 and in all ordinary chemical transforma- 
tions they yield substances of either equal or greater, but never less 
weight than their own. 

Many of these elements are familiar from our contact with 
them in every-day life; for example, oxygen, mercury, sulfur, 
iron, copper, lead, gold, silver, carbon, tin and zinc. A com- 
plete table of the elements may be found on p. 79. 

Compounds. These elements are capable of entering into 
combination with one another, thus giving rise to the almost 
innumerable substances known to chemists as chemical com- 
pounds. These chemical compounds show the following char- 
acteristics. They have entirely different properties from those 
of the elements from which they are formed. They are homo- 
geneous. The proportions by weight in which the elements 
combine is in accord with the law of constant proportions. The 
weight of the compound is always greater than that of any one 
of the elements from which it was formed and is equal to the 
sum of their weights. Probably the most important of all 
properties is that described by what is known as the law of the 
conservation of the elements. No ordinary chemical process (ex- 

1 The word "ordinary" is introduced here to exclude radio-active trans- 
formation, since it is possible that in these cases the above statement may 
not hold, 


eluding radio-active changes) ever produces from a compound any 
other elements than those which united for its formation, and these 
elements are reproduced in precisely the proportion by weight in 
which they combined for the formation of the compound. These 
points may be illustrated by mercuric oxide. It may be formed 
by the union of the two elements oxygen and mercury in definite 
proportion by weight. Its properties are very different from 
those of mercury and of oxygen, and no elements but mercury 
and oxygen have been obtained from it. With a proper under- 
standing of the meaning of the term " element," we may now 
resume the discussion of the element oxygen. 

Occurrence in Nature. As we have already seen, oxygen 
makes up about one-fifth of the volume of the air. That the air 
is a mixture of gases and not a chemical compound will be shown 
later. An element which is found in nature uncombined 
chemically with any other elements, although it may be mixed 
with them, is said to exist in the "free state." Since there is 
such a great quantity of air, oxygen in the free state is very 
plentiful. The compounds of oxygen are numerous and exceed- 
ingly abundant. More than 47 per cent, of the earth's crust and 
eight-ninths of the water is oxygen. 

Almost every one of the elements is capable of undergoing 
chemical change with oxygen. Each substance so formed 
differs radically in its properties from either oxygen or the other 
element. Its weight is equal to the sum of the weights of 
the other element and of the oxygen, and the two elements 
unite for the formation of the substance in perfectly definite pro- 
portion by weight. Therefore all substances formed in this way 
are chemical compounds, and since they all have one common 
component, oxygen, they are called oxides. We have, for ex- 
ample, oxide of mercury, oxide of iron, oxide of lead, etc. 

Preparation of Ozone. When oxygen is exposed to the in- 
fluence of the silent electrical discharge, (Fig. 6) it undergoes 
remarkable alterations in its properties, and is partially trans- 
formed into a new substance called ozone. By removing the 
latter as it is formed, the whole of the oxygen can be changed 
into ozone. 

Physical Properties of Ozone. When oxygen is converted into 
ozone, the volume of the ozone is only two-thirds that of the 


oxygen from which it was formed and consequently its density 
is one and one-half times that of oxygen or 0.002144. It boils at 
119C. and may be obtained as a deep blue liquid by passing 
ozonized oxygen through a tube cooled with liquid air. Gaseous 
ozone has a faint blue color and a very strong odor. It is some- 
what soluble in water. 


/.Several / *Tin Foil 'Inside of 'Glass Tube 

Thousand.. J 

1 ft/ft "Y 

FIG. 6. 

Chemical Properties. Gaseous ozone has a very irritating 
action on the .mucous membrane. A piece of bright silver when 
exposed for some time to ozone becomes blackened. A colorless 
solution of potassium iodide becomes dark brown from liberated 
iodine when exposed to the action of ozone although oxygen 
has no effect on it. 

If ozone be heated to a temperature of 250 to 300C. it is 
changed back to oxygen again with an increase in volume equal 
to the contraction which took place when the ozone was formed. 
The presence of metallic silver and also of several metallic oxides 
increases the rate at which this transformation of ozone into 
oxygen takes place at any given temperature. Many coloring 
substances are destroyed by ozone and hence it is used as a 
bleaching agent. It rapidly destroys the lower organisms, 
such as bacteria, and is therefore used as a disinfectant. In 
general its action is one of oxidation and differs from oxygen 
chiefly in that it is more vigorous and that usually one-third of 
the weight of ozone is used in oxidizing and two-thirds becomes 
ordinary oxygen. Liquid ozone is a very dangerous substance 
since it explodes with great violence either when brought in 
contact with an easily combustible substance or when its tem- 
perature has reached its boiling-point. The gas when compressed 
is also explosive. 

Uses for Ozone. As was mentioned before, ozone is used for 
bleaching of oils, paper, delicate fabrics, and flour, sterilization of 


water, and in addition is used for the purification of air and of 
starch, resinification of oils, and in the aging of liquors. Its use 
for the purification of air is of somewhat doubtful value, since 
to destroy bacteria in air the ozone must be strong enough to 
kill guinea pigs and in this strength would in all probability be 
dangerous to human beings. 

Allotropy. Oxygen may be converted into ozone or ozone 
into oxygen without any alteration in weight. From this the 
conclusion is drawn that no other substance combines with the 
oxygen for the formation of the ozone. This conclusion is con- 
firmed by the fact that when phosphorus is acted upon by ozone 
in one case and oxygen in another, the same substance, phos- 
phorus pentoxide, is formed in each case, and if equal weights of 
oxygen and ozone are taken, the weight of phosphorus pentoxide 
formed in each experiment will be the same. What then is the 
difference between oxygen and ozone? 

Whenever oxygen is transformed into ozone, energy is taken 
up; and when the ozone is changed back into oxygen, this energy 
reappears, generally in the form of heat. 

Ozone then is to be regarded as a modification of the element 
oxygen, possessing more energy than the latter. This property 
which oxygen has in common with several of the other elements 
of existing in two or more distinct forms, each having its own 
set of properties, even when the physical state of the different 
forms is the same, is called allotropy. 

Allotropy then is the property shown by certain elements of 
existing in more than one distinct form, each having the same 
physical state but possessing different properties. The different 
forms are spoken of as allotropic modifications of the element. 

The fundamental difference between the allotropic modifications 
of any one element consists of a difference in energy, although 
in some cases, and possibly in all, there is a difference in molar 
weight (see p. 68). 


Hydrogen was shown to be a distinct substance by Cavendish 
in 1766. It had been observed as an inflammable gas many 
years before, but was confused with other inflammable gases. 
That it is an element is shown by the fact that whenever it under- 
goes chemical changes the substances produced weigh more than 
the hydrogen which was transformed. Like oxygen it is a gas 
at ordinary temperatures. 

Occurrence. Hydrogen occurs free in nature in rather small 
quantities. It is found in the gases given off from some vol- 
canoes and fumaroles and is present in the air to the extent of 
something like one volume in 30,000 volumes of air. The quan- 
tity is so small that it is difficult to determine it with accuracy. 
In the combined state it is very abundant. Its chief compound 
is water, of which it forms 11.19 per cent, by weight, the re- 
mainder being oxygen. It occurs also in combination in coal, 
petroleum, natural gas, and in almost all substances of animal 
or vegetable origin. It is the essential element in a large and 
important class of substances known as the acids. 

Preparation of Hydrogen. Since water is so abundant and 
contains so much hydrogen it would' seem natural to suppose 
that hydrogen might be obtained from it economically by some 
such process as was used in the preparation of oxygen. 

The preparation of oxygen from its compounds is compara- 
tively simple because there are several compounds, for example, 
mercuric oxide and potassium chlorate, which will break up o'n 
being heated to a moderate temperature into gaseous oxygen 
and some other substance which is either liquid (mercury) or 
solid (potassium chloride) at ordinary temperatures, and can 
consequently be readily separated from the gaseous oxygen. 

In the case of hydrogen, neither water nor any other compound 
has these properties and, therefore, a different method of prepara- 
tion is required. At a fairly high temperature, water can be 




partially decomposed into hydrogen and oxygen, but there 
is no practical method for the mechanical separation of these 
two gases. However, if there were present in the heated steam 
something with which the oxygen in the water vapor could 
combine to form a compound which is either solid or liquid at 
ordinary temperature, it ought then to* be possible to separate 
this compound from the hydrogen and so get the latter in the 
pure state. 

Powdered Iron 
{'andiron Oxide 

FIG. 7. 

Metallic iron forms a compound with oxygen which is solid 
even at very high temperatures, and if steam be passed over 
finely divided iron, heated to redness in a tube (Fig. 7) iron 
oxide and hydrogen will be formed. The iron oxide will remain 
in the tube, while the hydrogen, mixed with much unchanged 
steam, will pass on over. The hydrogen may be collected by 
the displacement of water, as was done with oxygen, the steam 
which accompanies it being condensed in the water of the 
pneumatic trough. 

The place of the iron may be taken by several other metals. 
Magnesium, for example, may be used, in which case the tempera- 
ture may be very much lower. In fact magnesium will rapidly 
decompose boiHng water if a little of a magnesium salt is dissolved 
in the water to keep the surface of the magnesium free from the 
magnesium compound formed. 

Metallic calcium in the form of turnings decomposes cold 



water at a convenient rate, thus furnishing a ready means 
for the preparation of hydrogen on a small scale. 

The metals sodium and potassium act upon water so vigorously, 
even at ordinary temperatures, that only very small quantities 
of the metals can be brought in contact with water without 
giving rise to dangerous explosions. These methods consist 
essentially in treating the compound of hydrogen, water, with 
some other element which combines with the rest of the com- 
pound other than hydrogen and leaves the hydrogen free. 
These are examples of one of the most commonly used methods 
for the preparation of the elements from their compounds. 

In each of these cases hydrogen and a non-gaseous substance 
result from the action, and the non-gaseous substance contains 
the metal and the oxygen. In addition it sometimes contains 
some hydrogen as we shall see later. 

Preparation from Acids. As has been stated above, hydro- 
gen is present in all acids. As examples of acids may be men- 

Acid and 
Zinc -... 

FIG. 8. 

tioned sulfuric, hydrochloric, acetic acids, etc. When acids 
are mixed with water and brought in contact with many metals, 
hydrogen is given off, and each metal forms a compound with the 
rest of each acid. The hydrogen in each case comes from the acid 
and not from the water. Although most metals and acids will 
work in this way, in some cases the action may be violent, in 
others extremely slow. Metallic zinc with either dilute sulfuric 
or hydrochloric acid works very regularly and at a convenient 
rate, hence these substances are largely used in the laboratory 
for the preparation of hydrogen (Fig. 8). 


Hydrogen is similarly prepared on a commercial scale or for 
filling balloons by the action of sulfuric acid on iron, but the 
iron always contains some carbon, and 'the hydrogen obtained in 
this manner is contaminated with compounds of hydrogen and 

Electrolytic Preparation. When two pieces of platinum, a 
metal which is very resistant to chemical action, are placed in 
water containing a little sulfuric acid, and connected to a suitable 
source of electricity, a current of electricity passes through the 
dilute sulfuric acid from one piece of platinum to the other. At 
the same time, free oxygen appears at one of the platinum plates 
and free hydrogen at the other. The process is really quite com- 
plex, but since the water gradually disappears and hydrogen and 
oxygen appear in the proportion in which they will form water, 
while the quantity of sulfuric acid in the solution is unaltered, it 
amounts to the decomposition of water at the expense of electrical 
energy. Here water is decomposed with the absorption of 
electrical energy very much as mercuric oxide is decomposed at 
high temperature with the absorption of heat, which it will be 
remembered is a form of energy. Since the hydrogen and the 
oxygen are given off at different points, it is an easy matter to 
collect each separately. One of the methods for the commercial 
preparation of hydrogen is based upon this principle using, 
however, plates of lead instead of the very expensive platinum. 
A solution of sodium hydroxide may be used instead of the sul- 
furic acid, but in this case the metal plates should be of iron or 

Purification of the Hydrogen. The hydrogen prepared by the 
methods outlined above always contains water vapor, and usually 
compounds of hydrogen and carbon from the carbon contained 
in the metal. The hydrocarbons may be largely removed by 
passing the gas through a solution of potassium permanganate. 
To remove the water vapor, advantage is taken of the fact that 
there are several substances which vigorously retain water while 
having no action upon hydrogen. One of the most convenient 
of these is granulated calcium chloride. This is a white, very 
hygroscopic substance which is formed as a waste product in 
many chemical operations and is therefore cheap. The calcium 
chloride is usually placed in a tube through which the gas is 


passed. The column of calcium chloride should be as long as 
convenient and held in place by plugs of cotton wool (Fig. 8). 

Concentrated sulfuric acid is a far more efficient drying agent 
than calcium chloride and is largely used. Since this is a liquid 
it is either placed in a suitable container and the gas bubbled 
through it, or it is taken up by pumice stone and packed in dry- 
ing tubes or towers. 

Physical Properties. Hydrogen has no taste or color. The 
pure gas has no odor, but that prepared by the action of acids 
upon metals usually has a disagreeable odor due to the com- 
pounds of carbon and hydrogen which have been mentioned 
above. After the gas has been passed through a solution of 
potassium permanganate, it is so far purified that it loses the 
greater part if not all of the odor. 

Its density is smaller than that of any other known substance, 
being 0.00008986 grm. per cubic centimeter under standard con- 
ditions. This is perhaps the most important physical property 
of hydrogen, since upon it depends its use in balloons. The den- 
sity of pure dry air under standard conditions is 0.001293 grm. 
per cubic centimeter or 0.0012031 grm. per cubic centimeter 
more than the density of hydrogen, consequently 1 c.c. of 
hydrogen under standard conditions will have an upward flota- 
tion of 0.0012031 grm. when surrounded by air under the same 

The extreme levity of hydrogen may be easily shown by fill- 
ing a small rubber balloon with the gas when it will float in the 
air and exert a lifting power of something like a gram per liter. 
Or soap bubbles may be blown with the gas, and these upon being 
detached from the pipe will rise rapidly. 

Another way to demonstrate this same property is to fill two 
cylinders with hydrogen and hold one of them mouth upward and 
the other mouth downward. After a few minutes bring a flame 
to the mouth of each jar. It will be found that the one with the 
mouth upward contains only air, while the one with the mouth 
downward still contains hydrogen since the gas burns. 

Diffusion. Because of the very great difference in density 
between hydrogen and air one would expect that if two cylinders 
were brought together mouth to mouth (Fig. 9a) the upper one 
filled with hydrogen and the lower one with air, that the gases 




Cell filled 

would remain unmixed. But if they are left in this position for 
a few hours it will be found that both cylinders contain hydrogen 
and air. If the cylinders are of the same size and have been left 
in position for a sufficient length of time, the upper cylinder will 
contain as much air as the lower, and the lower will contain as 
much hydrogen as the upper. This 
spontaneous mixing of one gas with 
another, even against the action of 
gravity, is an example of the phenome- 
non called diffusion. 

All gases diffuse into one another and 
the process goes on until each gas is 
uniformly distributed through the entire 
space. If the pressure of each gas be 
measured separately, it will be found 
that diffusion continues until the pres- 
sure of any given gas is the same in all 
parts of the space; and further, it will 
be found that the pressure exerted by 
each gas is the same as that which the 
given mass of the gas would have exerted if it alone were 
occupying the entire space. This pressure is called its partial 
pressure. The total pressure of a gaseous mixture is the sum 
of all the partial pressures of its component gases. The volume 
of each gas is the total volume of the space occupied by the entire 
gaseous mixture. 

The law which describes these phenomena is known as Dalton's 
law of partial pressures, and may be stated as follows: Each 
gas, in a gaseous mixture, pervades the entire space occupied by the 
mixture and exerts a pressure equal to that which it would give if 
it were the only gas present; and the total pressure of the mixture 
is the sum of these partial pressures of the component gases'. 

This law may be expressed in symbols by representing the 
total volume by V, and that of the separate gases by vi, V2, 
v 3 . . . and the total pressure by P, and the partial pressures 
by Pi, pa, p 3 , ... then, 

FIG. 9. 


Vl = V2 = Va = V 
Pl + P2 + P3 = P 


While any gas will diffuse in^o any other gas no matter what 
the difference in density may be, there is often a great difference 
in the velocity with which the diffusion takes place. For 
example, in the case of the hydrogen and air cited above, the 
hydrogen diffuses downward more rapidly than the air diffuses 
upward. In fact in all cases the less dense gas will diffuse more 
rapidly than the denser. 

Something very nearly like this may be demonstrated quali- 
tatively by closing a porous cell (Fig. 96) such as is used in 
electric batteries, with a cork carrying a glass tube, and support- 
ing the whole arrangement in such a way that the glass tube dips 
into water. If a cylinder containing hydrogen be brought over 
the porous cell, bubbles will quickly begin to pass out from the 
glass tube. This action soon ceases and if the jar of hydrogen 
be removed the water will begin to ascend the tube. Both 
effects are due to the fact that the hydrogen being less dense will 
pass through the porous cell more rapidly than the air, produc- 
ing in the first case, some little pressure within the cell, and 
after the removal of the cylinder, the hydrogen passes out so 
much more rapidly than the air can enter that a partial vacuum 
results and the water rises. 

The relation between the velocity of diffusion and the density 
is that the velocities of diffusion vary inversely as the square 
root of the densities of the gases. 

Representing the velocities of diffusion by Ci and c 2 and the 
density by di and d2 this relation may be expressed by the 

Ci/c 2 

Hydrogen and the Gas Laws. The gas laws which were 
developed under oxygen describe very accurately the behavior 
of hydrogen under ordinary conditions; in fact, hydrogen follows 
them more closely than almost any other known gas. At very 
high pressures, however, it is found that hydrogen does not 
decrease in volume as rapidly as it ought according to Boyle's 
law. The discrepancy becomes greater as the pressure is in- 
creased. Other gases at ordinary temperatures are usually 
more compressible at a moderate pressure than they should be 
from Boyle's law, and then at higher pressures they behave 


like hydrogen and become less compressible. At higher tempera- 
tures many of them act like hydrogen and are less compressible 
from the start than they should be from Boyle's law. At 
very low temperatures hydrogen changes its behavior and is 
first more and then less compressible than it should be, thus 
behaving like the other gases. 

Liquid Hydrogen. Hydrogen cannot be liquefied by any 
pressure, however great, at ordinary temperatures, but must 
first be cooled to its critical temperature, 242C., when it will 
liquefy under a pressure of 13 atmospheres. Its boiling-point 
under atmospheric pressure is about 252. 5C. or only 20.5 A. 
By reducing the pressure to 5 cm. of mercury the temperature 
falls to about 14A. or 259C. At this temperature hydrogen 
is a transparent solid. Liquid hydrogen is colorless and has a 
density at its boiling-point of only 0.07. It is used to get very 
low temperatures, since substances placed in it will be very 
quickly cooled to the boiling-point of hydrogen. At this tem- 
perature all other liquids and all other gases except helium 
become solids. 

Solid Hydrogen. Solid hydrogen is a transparent ice-like 
substance with a melting-point of 258C. Its density is 0.076. 

Absorption of Hydrogen by Metals. Hydrogen is absorbed in 
appreciable quantities by many metals, but especially by the 
rare metal palladium. Under favorable conditions, palladium 
will take up 900 times its own volume of hydrogen, considerable 
heat being evolved during the process. This hydrogen will 
practically all be given up in a vacuum at ordinary temperatures 
or it may be driven off at atmospheric pressure by warming the 

The absorbed hydrogen does not enter into a chemical com- 
pound with the palladium but forms a solid solution. 

Spectrum of Hydrogen. When an electric discharge is passed 
through hydrogen under a pressure of a few tenths of a centimeter 
of mercury, the hydrogen glows and gives out light. This light 
appears rose red to the eye, but when it is examined with a 
spectroscope (Fig. 10), a very large number of lines are seen 
scattered all along from the red to the violet end of the spectrum. 
There are, however, four lines, red, green, blue and violet, which 
are much stronger than the others. Other gases under similar 




FIG. 10. 

conditions will give spectra but no other gas will give lines in 
precisely the same position in the spectrum as the hydrogen 
lines. These same lines will appear when hydrogen is caused to 
glow in any other way and their appearance is not prevented by 
the presence of other substances. Hence an examination of the 
spectrum is the surest means for the identification of hydro- 
gen. These same lines are found in the spectra of many of the 

stars, and we may safely say 
that hydrogen is present on 
them as well as upon the earth 
(see Frontispiece). 

The light from the sun and 
many of the stars shows dark 
lines at exactly the points 
where the hydrogen lines are 
bright. As will be explained, 
p. 293, these dark lines are due 

to hydrogen, and prove the presence of this element just as 
certainly as the bright lines. Therefore the spectroscope shows 
us that hydrogen is very widely distributed throughout our 
universe. In much the same way it has been established that 
a large number of the elements known to us here are also pres- 
ent in the sun and stars. 

Chemical Properties of Hydrogen. The most marked property 
of oxygen is that it will support combustion vigorously. Hydro- 
gen on the other hand will burn but not support the combustion 
of things which will burn in the air. This may be readily 
shown by bringing a lighted candle to the mouth of an inverted 
cylinder filled with hydrogen. The gas will take fire with a 
slight explosion and burn with an almost invisible flame at the 
mouth of the cylinder. Pass the candle up into the hydrogen 
and it will go out. Withdraw the candle slowly, and as the 
wick passes through the flame of the burning hydrogen it will 
be relighted. 

From what we have learned of combustion in studying oxygen, 
we would draw the conclusion, correctly, that hydrogen will 
combine with oxygen. It will also enter into combination with 
a large number of the other elements notably with chlorine, 
sulfur, nitrogen, and carbon. The conditions under which the 


combination will take place, and the substances formed will be 
discussed in connection with the several elements. 

A small stream of hydrogen issuing from a jet of platinum, or 
other difficultly fusible metal, burns quietly in the air with an in- 
tensely hot flame which is almost invisible. If the jet is made of 
glass the hydrogen flame will be colored yellow from the sodium of 
the glass. Water is produced by the burning of hydrogen, as may 
be shown by holding a glass cylinder over the flame. A dew at 
once appears on the glass which has all the properties of water and 
consequently is water. If all the water formed by the burning of 
a given quantity of hydrogen be collected and weighed, it will be 
found to weigh much more than the hydrogen itself, and therefore 
it follows that water is a compound of hydrogen with one or more 
of the constituents of the air. 

If the air be replaced by pure oxygen it will be found that the 
hydrogen will burn even more vigorously and water will be pro- 
duced as before. Hence we must conclude that water is a com- 
pound of hydrogen and oxygen and nothing else, a conclusion 
which is confirmed by the fact that the weight of the water 
formed is equal to the sum of the weights of the oxygen and 
hydrogen transformed. 

Hydrogen and oxygen may be mixed at ordinary temperatures 
and no perceptible combustion will take place, but if the mixture 
be brought in contact with a flame or heated at any point to a 
temperature of about 500 to 600C. a combination will take 
place very rapidly, and a flame will spread throughout the mix- 
ture almost instantaneously. The result is a rather powerful 
explosion due to the sudden expansion of the gases by the heat 
evolved during the rapid burning of the hydrogen. Because of 
its explosiveness, a mixture of hydrogen and oxygen in the pro- 
portion for the formation of water, i.e., two volumes of hydrogen 
to one volume of oxygen, is called detonating gas. This detonat- 
ing gas must be very carefully kept away from flames or in- 
candescent bodies. 

Because of the oxygen in the air a mixture of hydrogen and 
air in proper proportions is decidedly explosive. On this account 
hydrogen from a gas holder or from a generator should never be 
lighted until it has been tested and found to be pure. This may 
be done by collecting the escaping gas in an inverted test tube 


and bringing it in contact with a flame which is kept at a distance 
of several feet from the hydrogen apparatus. At first the gas 
in the tube will behave like air, but soon it will take fire, and the 
flame will rush up the test tube with a whistling noise. As the 
trials are repeated and the gas becomes more nearly pure hydro- 
gen, the explosions in the test tube first become stronger and then 
weaker until at last the hydrogen burns quietly at the mouth of 
the test tube for a sufficiently long time so that a flame may be 
carried back in this way from the burner to the hydrogen appara- 
tus and ignite the escaping gas. This method of lighting a jet 
of hydrogen is perfectly safe since so long as the gas in the appara- 
tus is explosive the explosions take place the instant the mouth 
of the test tube is brought to the flame, and no fire is left in the 
tube to ignite the jet upon the return of the tube. 

Oxy-hydrogen Blowpipe. Many important applications have 
been found for the very high temperature produced by hydrogen 
burning in oxygen. Because of the dangerously explosive char- 
acter of a mixture of these gases it is of course out of the ques- 
tion to mix them in one gas holder and burn the mixture. The 

gases must then be stored separately 
and burned from a special burner called 
the oxy-hydrogen blowpipe (Fig. 11) 
which is so arranged that the gases are 
conveyed to the burner in separate 
FIG < ~^~~ tubes and mixed just before they reach 

the orifice of the burner. In this way 
a flame may be produced whose temperature is something 
like 2,500C. In this flame silver boils, iron burns, and even 
platinum melts. If the flame be brought in contact with a 
piece of lime the latter becomes intensely heated and pro- 
duces an exceedingly bright light known as the calcium light, 
or lime light, and has often been used as a source of light for the 
stereopticon, but is now generally replaced by the electric arc 
or the nitrogen filled tungsten lamps. The oxy-hydrogen blow- 
pipe is largely used in the autogenous welding of metals, 
especially of lead. 

Formation of Water from Hydrogen and Oxygen Compounds. 
If hydrogen be passed over mercuric oxide at ordinary tem- 
peratures no perceptible action takes place, but if the oxide be 


gently warmed taking care that the temperature does not rise 
high enough to bring about its decomposition, water and metallic 
mercury will be rapidly formed. Hydrogen and iron oxide 
heated to a somewhat higher temperature will yield water and 
metallic iron. If tin oxide is used, metallic tin and water 
will be formed. Many other oxides will react in the same way 
with hydrogen so that we have here a rather general method for 
the laboratory preparation of a metal from its oxide. 

A very natural inference to draw from these facts is that 
the hydrogen has a stronger attraction for the oxygen than 
the metal has, and so robs the latter of its oxygen. This 
conception has often been used, and the attraction is called 
chemical affinity. That this conception is inadequate is shown 
by the following case. Iron oxide and hydrogen will react at a 
somewhat elevated temperature for the formation of water vapor 
and metallic iron. And the explanation of this in terms of the 
above conception would be that the hydrogen had a greater 
affinity for the oxygen than the iron had. But as was mentioned 
on p. 33, if water vapor be passed over metallic iron heated to 
the same temperature used in the above case, hydrogen and iron 
oxide will be formed. The explanation of this would be that the 
iron had a greater affinity for the oxygen than the hydrogen had 
and so jobbed the water of its oxygen. The action in the one 
case is just the reverse of that in the other and of course it. cannot 
be that the iron has both a greater and a less affinity for oxygen 
than hydrogen has. Nor can the reversal be due to change in 
affinity due to the difference in temperature because the reactions 
both take place at the same temperature. From this it can be 
seen at once that the explanation must be either abandoned or 

Such actions as these we have just been discussing are called re- 
versed or reversible reactions and are very common in chemistry. 
In fact they are more commonly met with than the irreversible 
reactions or those which go only in one direction. A careful 
study of these reactions is desirable because of their frequent 
occurrence, and also the ease with which they may be controlled 
when once understood. 

Probably the best way to obtain an understanding of re- 
versible reactions is to consider them as being composed of 


two simultaneous and opposite reactions, both taking place in 
the same vessel and each at its own rate. The rate of a re- 
action is measured by the mass of any one of the substances 
transformed per second and depends upon the nature of the 
substances undergoing the change, the temperature, and the 
mass of each of the reacting substances present per cubic centi- 
meter of reaction mixture. The mass of any given substance 
present per cubic centimeter is called its " active mass" or its 
concentration. 1 

At constant temperature, the rate at which any given substance 
reacts is greater, the higher the concentration of the reacting 

Let us choose for example the reaction between iron and water 
vapor and imagine that we seal up in a strong glass tube some 
water and some iron and raise the whole to a temperature of say 
400C. The water will react with the iron at a certain rate for 
the formation of hydrogen and iron oxide. As the water is used 
up the concentration of course decreases and the rate of reaction 
will diminish. On the other hand just as soon as some iron oxide 
and hydrogen have been formed they will begin to react for the 
re-formation of iron and water. The rate of this reaction will be 
zero at the start and gradually increase as the concentration of 
the hydrogen increases from its accumulation as the result of 
the first reaction. Hence the reaction between the iron and the 
water will go slower and slower as the water is used up and that 
between the iron oxide and the hydrogen will go faster and faster 
as the latter accumulates until finally just as much iron and 
water will be used up in the first reaction in a given time as is 
re-formed by the second reaction in the same time. When this 
state of affairs is reached, no further change in the concentra- 
tions will take place and the system is said to be in equilibrium. 
Equilibrium is defined as a condition which does not change with 
the passage of time. 

Catalysis. At temperatures of 500 to 600C., oxygen and 
hydrogen combine with explosive rapidity. As the temperature 
is lowered the rate rapidly decreases but is still appreciable at 

1 When dealing with a pure substance the concentration of that substance 
is the same as the density, but for solutions, the concentrations are different 
from the densities. 


250C., and there is no temperature at which we can say that the 
reaction just begins and below which it does not take place. We 
must conclude then that it is taking place at all temperatures 
and that even at ordinary temperatures it is still going on, 
though at an immeasurably slow rate. However, if one intro- 
duces into a mixture of hydrogen and oxygen a piece of platinum, 
the gases combine with appreciable rapidity at the surface of the 
platinum. Since in this combination a great amount of heat is 
developed, the temperature of the platinum is raised. This 
tends to increase the rate of combination and not infrequently 
the platinum is heated to incandescence and explodes the gases. 
It must be noticed that the rise in temperature of the platinum 
is the result of the combination of the gases and not the cause of 
the combination or combustion. What the platinum does in this 
case is to increase the rate of the reaction so that from being 
immeasurably small it goes on with visible rapidity. This 
effect takes place right at the surface of the platinum and, other 
things being equal, is proportional to the surface exposed. On 
this account, finely divided or spongy platinum acts more rapidly 
than massive platinum, and is almost certain to become so heated 
as to explode the gases. By mixing the spongy platinum in the 
proper proportion with clay and making the whole into pellets 
and burning them, the mass to be heated may be so increased 
that the temperature does not rise high 
enough to explode the gas mixture, and yet 
the gases will combine rapidly. 

An experiment to show this can be readily 
performed by fastening one of these pellets to 
a wire and supporting it several inches above 
the surface of water in a pan or other suit- 
able vessel (Fig. 12) and inverting a cylinder 
filled with a mixture of hydrogen and oxygen over the pellet so 
that the mouth of the cylinder is below the surface of the water. 
The gases quickly begin to combine as shown by the rise of 
the water into the cylinder. This continues until the pellet is 
covered with water or one or the other of the gases is used up, 
if one of the gases is in excess. 

If the platinum be examined after such an experiment it will 
be found to be unchanged. Here we have another instance of a 


substance merely by its presence increasing the rate at which a 
reaction takes place. (See case of manganese dioxide and 
potassium chlorate, p. 17.) Some other metals have the same 
property although to a smaller degree than platinum and it is 
very common to find that the rate of a reaction is greatly altered 
by the presence of a substance which is not changed by the 
reaction. We call actions of this kind catalytic actions. The 
substance through whose presence the rate of action is altered is 
called a catalyzer, and the whole process is spoken of as catalysis. 
Some catalyzers act as platinum does in this case and increase 
the rate of reaction ; these are called positive catalyzers. Others 
decrease the rate of reaction and are known as negative catalyzers. 
Catalyzers simply alter the rate at which a reaction is taking 
place and cannot bring about one which is not already going on. 
Cases of catalysis are of very frequent occurrence and will be 
often mentioned in the course of this work. 


Water is one of the most important as well as one* of the most 
widely distributed substances in nature. Some idea of its 
abundance may be gathered from the fact that the ocean con- 
tains about 300,000,000 cubic miles of water and each cubic 
mile weighs over 4,000,000,000 tons. Liquid water covers 
72 per cent, of the earth's surface; the atmosphere contains 
enormous quantities of water in the gaseous state, and the polar 
regions and the tops of lofty mountains are covered with solid 
water or ice. The crust of the earth is everywhere permeated 
by water, and all living organisms, both plants and animals, 

Block Tin 

FIG. 13. 

contain water as an absolutely essential portion of their struc- 
ture. The human body contains about 70 per cent, and the 
ordinary foods from 35 to 95 per cent, water. 

Water is a very good and general solvent and takes up more 
or less of nearly everything with which it comes in contact. 
Water is usually purified by the process of distillation which 
consists in heating the water until it boils and is converted into 




vapor and then, in a separate vessel (Fig. 13) , cooling the vapor 
and reconverting it into water. Most of the substances present 
in the impure water are much less easily vaporized than water and 
are left behind when the latter is boiled. Some impurities, such 
as ammonia and carbon dioxide, are more volatile than water and 
pass over with the first portions of the vapor. By rejecting the 
first part of the distillate, practically pure water may be obtained. 
Absolutely pure water cannot be prepared because it must come 
in contact with a vessel of some kind, and it will inevitably 
dissolve some of the latter, thus rendering the water impure. 
Stills and condensers of platinum or pure block tin are least 
attacked and give the purest water. 

For drinking purposes the dissolved impurities in natural 
fresh waters are not so important as the bacteria which they 
carry. Since the latter are in suspension and not in solution, 
^hey may be removed by very fine-pored filters, and in a sense 
the water is thereby purified. The same result may be obtained 
by killing the bacteria by ozone (see p. 30), or by bleaching 
powder (see p. 129), or by boiling. 

Physical Properties of Water. It is at ordinary temperatures 
a transparent, almost colorless liquid with a faint bluish tint 
which becomes apparent upon looking through a thick layer 
of water. This color can be seen in clear lakes and in the ocean. 

Most substances increase in volume more or less regularly 
with rising temperature, but water is a decided exception. 
If one starts from 0C. and warms the water its volume will 
decrease until a temperature of 4C. is reached, and then with 
rising temperature the volume increases. Since the density of 
a substance varies inversely as its volume, the temperature of 
maximum density of water is 4C. 

The following table gives some of the values for the density 
of water at different temperatures: 







40 C. 


4 C. 


60 C. 


10 C. 


80 C. 


20 C. 


100 C. 



Ice. At a temperature of 0C. water changes from the liquid 
to the solid state and is then known as ice. This change is on 
the border line between physical and chemical changes; so 
many of the properties of ice are different from those of water, 
as will be seen from what follows, that it may properly be 
considered a chemical change, and yet the transformation back 
and forth takes place so easily that many regard this as a 
physical change. 

Properties of Ice. Ice is a transparent solid of a faint bluish 
color which is seen only in large masses such as icebergs or 
glaciers. It is such a poor conductor of electricity that it 
becomes electrified when rubbed. Ice is almost twice as good 
a conductor of heat as water, but even then it is a very poor 

The density of ice at 0C. is 0.91674, which is much smaller 
than that of water at the same temperature which is 0.99987. 
This means that on freezing water increases nearly one-eleventh 
in volume, and also that ice will float on water. In these 
respects water is somewhat unusual, since most substances 
decrease in volume upon solidification, and the solid sinks in 
the liquid. 

In all climates where freezing and thawing take place, the 
great increase in volume when the water passes into ice has 
played a very important part in the disintegration of rocks. 
When the rain falls upon the rocks it fills even the small cracks. 
Upon freezing the increase in volume forces the crack to widen. 
When the ice melts the water runs down in the crack, and upon 
freezing spreads the pieces of rock still farther apart until 
presently the rock is split up under the repeated action of the 
freezing and thawing of the water. The bursting of water pipes 
when the water in them freezes, and the loosening of the soil , 
in the winter time are due to this same cause. 

Ice is a crystalline substance, and crystals may be defined as 
bodies many of whose properties vary with the direction in 
which they are measured. As a result of this variation a 
crystalline substance tends, when it takes on the solid state, 
to form bodies whose limiting surfaces meet at perfectly defi- 
nite angles, and so build up regular polyhedral forms. All other 
physical properties which can possibly do so vary with the 


direction in which they are measured. Some of these proper- 
ties are conductivity for heat and electricity, velocity of light, 
elasticity, etc. 

Those solids which are not crystalline are called amorphous 
substances, and glass is a good example of such a substance. All 
the properties of glass are the same in every direction. The 
property of forming crystals is very common among pure solid 
substances, while most amorphous substances are mixtures. 

Well-formed crystals often grow on the surface of water, 
but the best are found in the snow because they have been 
formed in the air with nothing to interfere. The snow crystals 
usually appear as six-pointed stars. 

Water in contact with ice cannot be cooled below 0C. nor 
warmed above this temperature, but so long as there is present 
an intimate mixture of ice and water the temperature is 0C. 
For this reason the ice point is generally used for a standard 

In the absence of ice, water may be cooled a number of degrees 
below zero without freezing. In this condition it is said to be 
supercooled. It is interesting to note that while water is cooling 
down from it continues to expand just as it does from 4 to 0. 
In fact there is no sudden change in the properties of water at 
except that at that temperature it becomes possible for ice to 
exist which it cannot do at higher temperatures. Water which 
has been cooled to a few degrees below zero is stable toward all 
kinds of changes except the introduction of ice. Let the very 
smallest piece of ice come in contact with the water and freezing 
at once starts and the temperature rises to 0C. 

If a test tube partially filled with ordinary distilled water be 
supported in a freezing mixture of ice and salt in such a manner 
that the surface of the water in the tube is a centimeter or so 
above the surface of the freezing mixture, the water will super- 
cool several degrees. By taking elaborate precautions water 
has been supercooled to 15C. 

Heat of Fusion. Whenever any liquid substance solidifies, 
heat is evolved; and conversely when a solid melts heat is ab- 
sorbed. The heat absorbed during the melting of 1 grm. of a 
solid without raising its temperature is called the heat of fusion 
of that solid. That evolved during the freezing of unit mass of 


the substance is called the heat of solidification. These heats are 
of equal numerical value and are both often called the heat of 
fusion. Ice behaves like all other solids in these respects and is 
only noteworthy because its heat of fusion is greater than that of 
any other substance except ammonia. 

Measurement of Heat. A quantity of heat is measured by 
determining the change in temperature which it will produce in a 
known weight of water; the heat capacity of the water, i.e., 
the number of joules required to raise the temperature of 1 
grm. of water 1, having been carefully determined. The heat 
capacity is also known as the specific heat and varies somewhat 
with the temperature. Another unit is that quantity of heat 
which raises the temperature of 1 grm. of water 1C. This is 
called the small calorie (abbreviated "cal.") Since the specific 
heat of water varies with the temperature, it is necessary to 
specify the temperature of the water. A larger and very generally 
used unit is the large or kilogram calorie. It is the heat necessary 
to rise the temperature of 1 kgm. of water 1C. It is abbreviated 
Cal. and equals 1,000 small calories. 

The heat of fusion of ice, that is, the heat required to melt 1 
grm. of ice, is 334 joules or 80 cal. If, therefore, 80 grm. of water 
at 1 be mixed with 1 grm. of ice, 81 grm. of water at will be 

In technical work in England and the United States, another 
unit for the measurement of heat is in use. It is known as the 
British thermal unit and is abbreviated B.t.u. It is the 
quantity of heat required to raise the temperature of 1 Ib. of 
water 1F. 

The relation between the small calorie and the erg has been 
carefully worked out by Joule and others, and a small calorie at 
15 is equal to 41,880,000 ergs or 4.188 joules. 

The fact that the temperature of a mixture of ice and pure 
water is always finds its explanation in this heat of fusion. If 
heat is applied to the mixture, instead of raising the temperature 
some of the ice melts and absorbs the heat. If heat is taken 
away from the mixture, instead of the temperature's falling 
some of the water freezes and supplies the heat. Because of the 
constancy of this temperature ice finds many applications in 
science. Because of its cheapness, its great heat of fusion, and 



its rather low melting-point, enormous quantities of ice are used 
in every day life for the preservation of food since the destruc- 
tive changes take place much more slowly in the food cooled 
by ice than at ordinary temperatures. They are like all other 
chemical changes in that they are slowed down by lowering the 

When ice and common salt are brought together a solution 
results, the freezing-point of which is much lower than 0C. 
This is the basis of the use of ice and salt for freezing ice cream, 
etc. By using three parts of ice to one of salt a temperature of 
18C. can be obtained; this was taken by Fahrenheit as the 
zero point for his thermometer. The 
common salt may be replaced by other 
soluble substances because the freezing- 
point of any solution is lower than that 
Vacuum-^ o f the pure liquid. .Most substances 

will not give as low a temperature as 
salt, but a few will give even lower. 
However, because of its cheapness and 
convenience salt is the best to use. The 
use of salt for the removal of ice and 
snow from sidewalks, etc., depends upon 
this same phenomenon. 

Gaseous Water. Both water and ice 
are very easily transformed into gaseous 
water which is commonly known as 
k water vapor at ordinary and as steam at 
higher temperatures, although there is 
no real difference between them. To 
bring about this transformation, all that 
is necessary is to introduce the water or 
FlG - 14 - ice into a vessel which it does not 

entirely fill. Under these circumstances 

the formation of the gaseous water will begin at once and the 
vapor will soon fill all the space unoccupied by the other 
modification. The transformation of water or ice into water 
vapor is called evaporation. Water vapor is a gaseous sub- 
stance and its behavior is fairly well described by the gas 
laws. This is particularly the case if the pressure is low. 





Like all other gaseous substances, it exerts a pressure upon 
the walls of the containing vessel, and experiment shows 
that when water is brought into a space which it does not fill, 
evaporation takes place until either all the water has been 
converted into vapor or the pressure of the water vapor has 
reached a certain definite limit. This limit varies with the 
temperature but is always the same at any given temperature, 
and is entirely independent of the shape or size of the vessel 
or the relative volume of the liquid or the solid and the water 
vapor. We are dealing here then with a state of equilibrium 
between the water vapor and the liquid water or the ice. The 
limiting pressure at which the water vapor is in equilibrium 
with the other modification is called the vapor pressure of the 
water, or of ice as the case may be. It may be measured in the 
apparatus shown in Fig. 14. The vertical distance from a to b 
measures the vapor pressure. 

The following table gives the vapor pressure of water at 
different temperatures: 


ture C 

in cm. 


in cm. 


in cm. 










80 . 










28 . 


































From Dal ton's law of partial pressures and the fact that water 
vapor is gaseous, it can be readily seen that the vapor pressure of 
water will be the same whether the evaporation takes place in a 
vacuum or into a space filled with gases which do not dissolve 
to any marked degree in water. If the gases are very soluble 
they will alter the water, and hence the vapor pressure of the 


When heat is applied to water in an open vessel, the tempera- 
ture of the water gradually rises and finally small bubbles of 
gas appear in the liquid, rise to the surface and burst. These 
small bubbles consist of the gases dissolved in the water and of 
water vapor. The temperature continues to rise and these 
small bubbles are replaced by larger ones which consist of 
water vapor alone. This phenomenon is then called boiling. 
After the water has begun to boil, the temperature does not 

A little reflection will convince one that a bubble of vapor 
down in the body of the liquid must be under a pressure equal 
to that of the atmosphere plus that due to the column of liquid 
above it, and that due to the action of surface tension. 

The vapor pressure of the boiling liquid must be equal to these. 
Neglecting the last two factors which are usually small, we may 
say that the boiling-point of a liquid is the temperature at which 
the vapor pressure of the liquid is equal to the atmospheric pressure. 
Since the boiling-point is the temperature at which the vapor 
pressure of the liquid is equal to the external pressure, the boiling- 
point will vary as the pressure changes, and will rise as the pres- 
sure is increased. The boiling-point of water under a pressure 
of one atmosphere is 100C.; at two atmospheres the boiling- 
point is 121C.; under ten atmospheres it is 180C., and con- 
tinues to rise up to the critical point when the boiling-point of 
water is 374 C. under a pressure of 200 atmospheres. By lower- 
ing the pressure, water may be made to boil at temperatures 
below 100C. With the aid of a good water vacuum pump, 
water may be made to boil at temperatures below 30C. If the 
pressure be reduced to less than 0.4 cm. of mercury, the boiling- 
point will fall below the freezing-point, and the water will freeze 
and boil at the same time. 

Use is often made in the laboratory of the fact that the tem- 
perature of water boiling in an open vessel, and also the steam 
from it is never far from 100C. For example it is often necessary 
to heat a substance to about 100C. and at the same time abso- 
lutely essential that the temperature shall not go higher. This 
can be easily done by heating it on what is known as a water 
bath. The simplest form of a water bath consists of a beaker 
half full of water and loosely covered with an evaporating dish 


or other suitable vessel containing the substance to be heated. 
If now the water be boiled, the steam will come in contact with 
the dish and quickly raise the temperature to about 100C. 
but never above the boiling-point of water at the pressure of the 

Humidity of the Air. From the very wide distribution of 
liquid water upon the earth's surface, it of course follows that 
water vapor is always present in the air. One might expect that 
the partial pressure of water vapor in the air would be equal to 
the vapor pressure of the water. However, this is rarely the 
case because the temperature of a body of water is usually lower 
than that of the surrounding land area, and even if the partial 
pressure of the water vapor in the air over the water becomes as 
great as the vapor pressure of the water at the temperature of 
the water, as soon as the air moves over the land, its temperature 
will rise and the partial pressure of the water vapor will then be 
less than the vapor pressure of the water at the temperature of 
the air. This is sometimes expressed by saying that the air is 
only partially saturated, carrying the idea that the air dissolves 
the water much as water dissolves salt. The notion is only 
partly correct. The term "relative humidity" is often used in 
this connection and is the ratio of the partial pressure of the 
water vapor in the air to the vapor pressure of water at the 
temperature of the air. 

The water vapor in the air is of great consequence in chemistry 
because of the fact that everything which is exposed to the air 
takes on more or less water. 'Some things take on so much water 
that they pass into a liquid solution or deliquesce, while others 
simply condense a thin film of water upon their surface. This 
last is the behavior of glass and other substances that do not 
dissolve. This water condensed upon the surface often gives 
a great deal of trouble in making accurate weighings, because of 
the fact that it will vary markedly with the relative humidity 
and temperature of the air. The most efficient way to remove 
this film of liquid water is to heat the object and drive the water 
off as vapor. However, it must be noticed that a film of water 
does not behave exactly like liquid water in that its vapor pres- 
sure is very much smaller than that of liquid water at the same 
temperature, and consequently the temperature has to be very 


high to remove the water thoroughly. In some cases it is not 
possible to heat the object to drive off the water, and in this case 
it must be placed in a closed vessel called a desiccator which 
contains one of the drying agents already mentioned. Under 
these conditions, the water will vaporize from the object and 
slowly diffuse to the desiccating substance. The process is a 
very slow one and will be greatly hastened by removing the air, 
since the latter interferes with the diffusion of the vapor. 

Heat of Vaporization of Water. Whenever any liquid, water 
included, passes into the gaseous state heat is absorbed unless the 
transformation takes place at the critical point. This absorption 
of heat is not accompanied by a rise in temperature and simply 
represents the difference in energy between the liquid and the 
gaseous substance and the work done in the transformation. 
This absorption of heat is often ascribed to the work done in 
separating the particles of water from each other, since gaseous 
water occupies a very much larger volume than the liquid water 
from which it was formed. However, this cannot be the explana- 
tion because there is an increase in volume when water freezes, 
and yet heat is given out in this case. The heat absorbed in the 
transformation of 1 grm. of a liquid into its vapor, without rise 
in temperature, is called its heat of vaporization. In the case of 
water it is 537 cal. or 2,245 joules at 100C. or 606.5 cal. or 2,535 
joules at 0C. So as the temperature rises, the heat of vaporiza- 
tion decreases and becomes zero at the critical temperature. 

Steam is extensively used for the heating of buildings, because 
such large quantities of heat may be transmitted with the 
movement of only a small amount of substance. 

Effect of Pressure on the Melting-point of Ice. Ice and water 
exist in equilibrium over a certain range of temperature and 
pressure, for each temperature there is a perfectly definite 
pressure of equilibrium. Experiment has shown that the 
melting-point of ice is lowered by 0.0072 for each atmosphere by 
which the pressure is increased. An increase in the pressure of 
139 atmospheres is required to lower the temperature 1C. A 
mixture of ice and water then, can have a temperature of 0C. 
only when under a certain pressure. By definition of 0C. this 
is the standard pressure. A change in the barometric pressure 
should then affect the melting-point of ice and the alteration 


should amount to about 0.0001C. for a change in the barometric 
pressure of 1 cm. 

Although the lowering of the temperature for each atmos- 
phere increase of pressure is very small, it is possible by pressure 
alone to keep water in the liquid state at 24C. The pressure 
necessary for this may be most readily obtained by freezing a 
part of the water contained in a very strong steel vessel. The 
freezing of only a small part of the water will produce an enor- 
mous pressure because of the marked increase in the volume of 
water when it is changed into ice. 

The Law of Mobile Equilibrium. The melting-point of prac- 
tically every solid is altered to some extent by a change in pres- 
sure. An increase in pressure lowers the melting-point of some 
few but raises that of most substances. In every case an in- 
crease in pressure causes some of the less dense form of the 
substance to change into some of the denser form, and so pro- 
duces a contraction which tends to relieve the pressure. To 
maintain a balance under such conditions between the relative 
quantities of the two forms, the temperature must be altered 
in such a direction as will of itself tend to produce more of 
the less dense form. Usually the solid is denser than the 
liquid, and the tendency is for it to form when the pressure is 
increased. So the melting-point of most substances is raised. 
The case is just the other way around with water for the ice 
is less dense than the water and the melting-point is lowered. 
In every instance we are dealing with an equilibrium between 
a solid and a liquid, and the shifting of this equilibrium under 
the action of a constraint in the form of pressure. The equilib- 
rium is always shifted in such a way that the pressure tends 
to be relieved. These are special cases of a very general law 
which applies to all kinds of equilibria both physical and chem- 
ical. This law may be stated as follows: // a constraint is put 
upon a system in equilibrium whereby the equilibrium is altered, 
that reaction will follow which tends to decrease the constraint. 

Solutions. It is a matter of every-day experience that many 
substances will dissolve in water. We call these water or 
aqueous solutions of the substances. A solution, it will be 
recalled, is a homogeneous mixture of two or more substances, 
the relative proportions of which may vary continuously within 


certain limits and whose properties vary with the proportions of 
the components. Solutions may exist in all three of the physical 
states, and it is simply the circumstance that we are so familiar 
with water solutions that makes us tend to think of a solution 
as a liquid. 

In connection with solutions two terms, solute and solvent, are 
in constant use, and their meaning should be understood. The 
substance which is dissolved is called the solute, while that in 
which the solute is dissolved is called the solvent. A solution 
of salt in water is obtained by dissolving the solute salt in the 
solvent water. The solute may be solid, liquid, or gaseous, and 
so may the solvent. As an example of a liquid dissolving in a 
gas, we may take the vaporization of water into air. 

Aqueous solutions are of great importance in chemistry because 
of the fact that many chemical processes take place very readily 
in them; one reason for this being, apparently, that the water 
will bring together into one liquid, solid, liquid or gaseous solutes 
in the very most intimate mixture. 

In general, solutions have properties which are a sort of mean 
between the properties of the two components, so these aqueous 
solutions will have in the main the properties of water. One 
very important circumstance to note about solutions is that a 
solute always lowers the vapor pressure of its solvent upon pass- 
ing into solution. From this it follows that if the solute is non- 
volatile, the vapor pressure of the solution will always be less at 
any given temperature than the vapor pressure of the pure 
solvent. In case the solute is volatile, the partial pressure of 
the solvent is lowered as before, but the vapor pressure, which 
is the sum of the partial pressures of the solvent and of the 
solute, may or may not be lower than that of the solvent in 
the pure state. For solutions which are dilute, the lowering of the 
vapor pressure is proportional to the mass of the solute dissolved 
in a given fixed mass of the solvent, at constant temperature. 

In Fig. 15, the curve AOB is the vapor pressure curve for the 
pure solvent, and the curve CO'D that for the solution. It will 
be noticed that the curve for the solution lies below that for the 
solvent, and that they are not parallel. The curve EO'O is the 
vapor pressure curve for the ice. The line MN represents the 
atmospheric pressure. From the definition of boiling-point it 


will be seen that the temperature corresponding to the points 
B and D represent the boiling-points of the pure solvent and of 
the solution respectively. This also shows that owing to the 
decrease in the vapor pressure caused by the solute, the boiling- 
point of the solution is higher than that of the solvent. This is 
true in every case where the solute is non-volatile. If the solute 
is volatile, the boiling-point of the solution may be higher or 
lower than that of the pure solvent depending upon whether the 
partial pressure of the solute is less than or greater than the 
lowering of the vapor pressure of the solvent. 


M B D N 

FIG. 15. 

The freezing-point of all solutions from which the solvent 
separates as a pure solid is lower than that of the pure solvent. 
In order to understand this, it will be necessary to consider the 
relations between the vapor pressure of the solid and of the 
liquid. It seems very natural to assume that ice has no vapor 
pressure, but this is a mistake for ice has a measurable vapor 
pressure even when cooled to many degrees below zero. This 
pressure is 0.46 cm. at zero degrees, and becomes smaller as the 
temperature falls. As may be seen by reference to Fig. 15, 
the vapor pressure curve for ice is not a continuation of that 
for water, but lies below the water curve except at point 
O, where the curves meet. It is only at this point, O, then 
that the solid and liquid have the same vapor pressure and 
can be in equilibrium with each other. At temperatures below 


this point the vapor pressure of the water is higher than that of 
the ice and a little thought will convince anyone that the two 
could not continue to exist side by side because the water would 
evaporate and condense on the ice as the solid, and this would 
go on until all the water had disappeared. A little reflection 
will show that a system must be in equilibrium in every way to 
be in equilibrium in any way. Therefore it follows that since 
at the freezing-point the solid and the liquid are in equilibrium, 
the vapor pressure of the two must be the same. We may then 
frame a definition of the freezing-point so that it shall be similar 
to that for the boiling-point. The freezing-point of a substance 
is the temperature at which the vapor pressure of the solid is the 
same as that of the liquid modification. In Fig. 15 this is at the 
point O, where the two curves AOB and EO'O cut. Correspond- 
ingly the freezing-point of the solution will be the temperature 
at which the vapor pressure of the solution equals that of the ice, 1 
and this will be at the point, 0', on the curves and will always 
be at a lower temperature than that of the pure solvent. From 
these considerations it will be seen that the solute lowers the 
vapor pressure of the solvent and that this results in the solu- 
tions boiling at a higher, and freezing at a lower temperature 
than the pure solvent. No real explanation for the lowering 
of the vapor pressure can be given, but we can picture it to our- 
selves by assuming that things dissolve because they have an 
attraction for each other, and that, because of this attraction the 
solvent finds it more difficult to pass over into the gaseous state 
from a solution than when it is pure. 

When a solid salt is brought into contact with water, some of 
it dissolves. If the salt is difficultly soluble, like silver chloride, 
the weight of the salt which a given weight of water, say 100 
grm., can take up is very small indeed; but in the case of an 
easily soluble compound like calcium chloride, the mass which 
can be dissolved by the same weight of water is very large. 
However, there is at each temperature and pressure a definite 
limit to the weight of a salt which a given weight of water will 
dissolve. A solution in which this limit has been reached, and 

1 The student should remember that when a solution freezes the solid 
which separates almost always consists of crystals of the pure solvent and 
not of the solution. 


which is in equilibrium with some of the solid solute, is called 
a saturated solution. The ratio between the weights of the 
solute and the solvent in a saturated solution is called the solu- 
bility of the solute. It is usually expressed in terms of parts by 
weight of solute in 100 parts by weight of solvent. 

As a rule the solubility of a salt is markedly altered by changes 
of temperature. Generally, but by no means always, the solu- 
bility increases with rising temperature. The solubility is very 
much less sensitive to changes in pressure than to changes in 
temperature, in fact it requires a change in the pressure of 
hundreds of atmospheres to make a measurable change in the 
solubility. So small is the influence of pressure that for all 
practical purposes it may be neglected. 

If less of the salt be added to the solvent than it is capable of 
dissolving, all will be taken up. The solution is said to be 
unsaturated because it can still take up more of the salt. If more 
of the salt be added than the solvent can dissolve, that quantity 
which corresponds to the solubility will pass into solution and 
the excess will be left undissolved. The concentration of the 
salt in solution is absolutely independent of whether this excess 
be large or small. 

Just as a liquid may be supercooled if the solid modification be 
excluded, so a solution may be prepared which contains more of 
the solute than corresponds to its solubility if the solid solute is 
not present. Such solutions are called supersaturated. In the 
absence of the solid solute, they are stable toward most changes ; 
but when brought in contact with even a very small piece of the 
solid, all the excess of solute over that corresponding to its 
solubility quickly separates out in the solid form. 

Application of the Law of Mobile Equilibrium. We must 
go to the law of mobile equilibrium for information as to the 
direction of the change of solubility for a given change in tem- 
perature or pressure. 

This law says that if a system in equilibrium is subjected to a 
constraint whereby the equilibrium is altered, that reaction will 
follow which tends to relieve the contraint. If the constraint 
consists in the addition of heat whereby the temperature of the 
system is raised, the change which will follow will be accompanied 
by the absorption of heat. If a salt dissolves in its almost 


saturated solution with the absorption of heat, more salt will 
pass into solution when the system is heated. In other words, 
the solubility of the salt increases with rising temperature. On 
the other hand, if a salt dissolves in its almost saturated solution 
with the evolution of heat, it will pass out of solution into the 
solid state with the absorption of heat. So if such a solution be 
heated, solid salt will separate as the temperature rises, or the 
solubility will decrease with rising temperature. A salt which 
dissolves without any heat effect will not change its solubility 
with the temperature. This is very nearly true of sodium chlo- 
ride. The heat of solution of a salt in its saturated solution is 
often very different, not only in magnitude but also in sign, from 
its heat of solution in the pure solvent. It is only the former 
heat effect which is of significance in this connection. 

If the pressure be increased, that change will take place which 
is accompanied by a decrease in volume, since this will tend to 
relieve the pressure. The solubility of the salt will increase with 
the pressure in case the more concentrated solution has a smaller 
volume than the saturated solution together with the solid salt. 
But if the volume of the system would be decreased by the 
separation of the solid salt, the solubility of the salt will decrease 
with increasing pressure. 

The Chemical Properties of Water. Water can react chem- 
ically in essentially two different ways, that is, it may combine di- 
rectly, in which case only one new substance is formed, or it 
may react because of its being a compound of hydrogen and 
oxjrgen. In this case two or more new substances are formed. 
As examples of direct combination, we might cite the union 
of water with lime (calcium oxide) to form calcium hydroxide; 
with anhydrous copper sulfate to form crystallized copper 
sulfate or blue vitriol; with anhydrous calcium chloride to form 
crystallized calcium chloride. These compounds of water and 
the salts such as the copper sulfate and the calcium chloride 
are called hydrates or salts with water of crystallization. The 
hydrates differ greatly in their stability, and some are so easily 
decomposed into water and the rest of the compound that many 
have been led to believe that they are essentially different from 
Other chemical compounds, and that the water exists in them 
in some peculiar way as water. They have, however, all the 

WA TER 63 

characteristics of chemical compounds, that is, follow the law 
of definite proportions, differ radically in their properties from 
their components, etc., so there is no reason for setting up a 
special class for these substances. 

As cases of the behavior of water as a compound of hydrogen 
and oxygen, we may call to mind the reactions between water 
and iron, or magnesium, calcium or sodium whereby hydrogen 
and a compound of the metal with oxygen or with oxygen and 
hydrogen are formed. It is important to notice in this connection 
that it is much more difficult to decompose water into its elements 
than it is such a compound as mercuric oxide; on this account 
water is called a stable substance. 

The Reaction of Sodium and Water. The action of sodium 
upon water is so very vigorous that only very small quantities 
of sodium, a piece not larger than a pea, should be used at a 
time. Larger quantities are very likely to produce explosions, 
from the very great volumes of gas and steam liberated, which 
are as violent as those due to an equal weight of dynamite. 
When sodium acts upon water, hydrogen and a solution of sodium 
hydroxide are formed. Sodium hydroxide is a white easily 
soluble solid which is a highly caustic alkali and hence is often 
called caustic soda. When sodium hydroxide is exposed to the 
air, it becomes moist and finally passes into a liquid, which is 
identical in its properties with a strong water solution of the 
substance and hence is such a solution. The water necessary 
for the formation of this solution comes from the air. Sodium 
hydroxide and things like it which take up water vapor from 
the air and form solutions are called deliquescent substances. 
All such subtances are very freely soluble and have the common 
characteristic that their saturated solutions give a smaller 
vapor pressure of water than the partial pressure of the water 
vapor in the air. If this were not so the substance could not 
deliquesce. Many of the salts containing water of crystalli- 
zation decompose readily into their components and give an 
easily measurable vapor pressure of water. If the partial 
pressure of the water vapor in the air is less than the vapor pres- 
sure of the water from the salt, the latter will decompose and 
the crystals will fall to powder. The substance is then said 


to effloresce. At any given temperature the vapor pressure of a 
saturated solution of a substance or that of water from a hydrate 
is constant, but the partial pressure of the water vapor in the 
air varies from time to time, so a substance may deliquesce 
one day, be stable the next, and even effloresce on the day 
after; its behavior being determined by the partial pressure 
of the water vapor in the air and the vapor pressure of the water 
from the substance. 

In ordinary weather calcium chloride will deliquesce, and 
even after the salt has all dissolved, the solution will continue 
to take up water from the air until the partial pressure of the 
water vapor in the air is equal to the vapor pressure of the 
solution. If such a solution be allowed to stand in the labora- 
tory until winter comes on, it will be found that the water 
will gradually evaporate and finally crystals of calcium chloride 
with water of crystallization will form. In the depth of winter 
these crystals sometimes fall to a powder or effloresce. So 
calcium chloride is both a deliquescent and an efflorescent sub- 
stance. The explanation being of course that in the summer 
the partial pressure of the water vapor in the air is high, while 
in the winter it is low indoors on account of the low temperatures 
prevailing outside. 

The Composition of Water. Water is, of course, a chemical 
compound of hydrogen and oxygen and hence from the law of 
definite proportions the relative proportion by weight in which 
these elements combine to form water will be fixed and invariable. 
This proportion is a very important property, not only of hydro- 
gen and oxygen, but also of water. 

There are two general methods of procedure for determining 
the composition of a substance. We may decompose the sub- 
stance into its elements and determine the proportion by weight 
in which each is present. This is called analysis. Or we may 
bring together the elements which go to make up the substance 
and make the conditions such that they will combine to form the 
desired compound. By carefully noting just what weight of 
each element disappears in the formation of a given weight of the 
compound, the quantitative composition of the substance may 
be determined. This is called synthesis. 

For the analysis of water, advantage may be taken of the fact 



mentioned under the discussion of the methods for the prepara- 
tion of hydrogen, that when a current of electricity is passed 
through water between two plates of platinum, hydrogen ap- 
pears at one of the platinum plates and oxygen at the other. 
In this process the platinum plates are unaffected, and the 
weight of the hydrogen plus the weight of the oxygen equals 
the weight of the water which disappears, thus showing that 
these substances are the only ones formed in the process. By 
collecting the gases separately and determin- 
ing their weights, a quantitative analysis of 
the water may be made. Since the densities 
of the gases are known, it is not necessary 
to weigh them; all that is required is that 
their volumes be measured under known 
conditions of temperature and pressure 
(Fig. 16), and from these data the volumes 
under standard conditions may be readily 
calculated. The volumes under standard 
conditions times the densities of the gases 
will, of course, give the weights of the gases, 
and therefore the proportion by weight in 
which they combine to form water which 
is the result desired. 

When this operation is performed it turns 
out that if the gases be measured under 
the same conditions of temperature and 
perssure, that the volume of hydrogen is 
almost exactly twice that of the oxygen. 
The ratio according to the best determina- 
tions being 2.0025 volumes of hydrogen to 
1.0000 volume of oxygen. The density of 
hydrogen is 0.00008986 and of oxygen 0.0014293 and therefore 
the ratio by weight is 2.0025 X 0.00008986 divided by 1.0000 X 
0.0014293, or 1.000 of hydrogen to 7.943 of oxygen or roughly 
1 to 8. The percentage composition of water is then 11.19 per 
cent, hydrogen and 88.81 per cent, oxygen. 

One of the standard methods which has been much used for the 
quantitative synthesis of water is based upon the reaction taking 
place between copper oxide and hydrogen at a somewhat elevated 

FIG. 16. 



temperature whereby water and copper are produced. The 
process may be briefly outlined as follows: Pure dry hydrogen 
(Fig. 17), is passed over a weighed quantity of copper oxide at 
a little below the temperature of redness. The water formed 
is collected by absorption in a weighed quantity of calcium 
chloride. The increase in weight of the calcium chloride at the 
end of the experiment over that at the beginning gives the 
weight of the water formed. The difference between the weight 
of the copper oxide and of the copper formed from it gives the 
weight of the oxygen which was transformed into the water. 
The difference between the weight of the water and that of the 
oxygen is, of course, the weight of the hydrogen. The composi- 
tion of water from data obtained in this way checks very closely 
with that given above. 



FIG. 17. 

Another method for the quantitative synthesis of water con- 
sists in determining the proportion by volume in which hydro- 
gen and oxygen unite to form water. From this ratio and the 
densities of the gases, the proportion by weight may be easily 
obtained as was shown above. Since a mixture of hydrogen 
and oxygen may be caused to explode by passing an electric 
spark through the mixture at any point this synthesis is very 
easily carried out. The apparatus necessary consists of a 
thick- walled glass vessel (Fig. 18), strong enough to stand the 
explosion of the gases and having two platinum wires sealed 
through its walls. Carefully measured volumes of hydrogen 
and of oxygen are introduced into the vessel, mixed and ex- 
ploded. The residual gas is then measured and examined 



to see whether it is hydrogen or oxygen. The results show as 
before that the gases combine in the proportion of very nearly 
two volumes of hydrogen to one of oxygen. If the gases are 
mixed in just this proportion no residue is left; but if an 
excess of either one is used, the excess of that gas is left un- 
changed. If this synthesis is carried out at temperatures above 
the boiling-point of water and at pressures somewhat below that 
of the atmosphere so that the water formed shall stay in the 
gaseous state, as water vapor, the volume of this water vapor 
is very nearly the same as that of the 
hydrogen from which it was formed, both 
measured at the same temperature and 
pressure. So the following very simple 
relation exists between the volumes of the 
gaseous substances involved in this reac- 
tion: two parts by volume of hydrogen 
will combine with one part by volume of 
oxygen to form two parts by volume of 
water vapor, all being measured under 
the same conditions of temperature and 

When other gases enter into reaction 
with one another, relations between the 
combining volumes are always very 
simple, often even simpler than in the 
case given above. This behavior of 
gases is described by the law known as 
Gay Lussac's law of the combining vol- 
umes of gases. Whenever in any chem- 
ical reaction two or more gases appear or 

disappear , they do so in the ratio of small whole numbers by volume. 
In the case given above, the ratios are 2:1:2. 

When we come to compare the proportions by volume in 
which solids or liquids react, we find no suggestion of the exist- 
ence of a simple relation, nor is there any simple proportion by 
weight in which substances react. We find here another instance 
of the great simplicity of the behavior of gases. 

It will be recalled that Boyle's and Gay Lussac's laws do not 
exactly describe the behavior of any gas, but that they all vary 



somewhat from these. Similarly Gay Lussac's law of combining 
volumes is not an exact law as may be seen from the fact that 
hydrogen and oxygen do not combine in the ratio of 2 : 1 as called 
for by the law, but in the ratio of 2.0025 : 1. The deviations, 
however, are not large, in most cases, and the law is very useful 
and important. 

Molar Weight. The conception of molar weights which is 
about to be developed is highly useful since with its aid the inter- 
pretation of many apparently complicated chemical phenomena 
becomes simple. 

The molar weight of a gaseous substance is conceived of as 
being proportional to the density of the gas. For reasons 
which cannot be brought out until later, the molar weight of 
oxygen is taken as 32. At the proper point the reasons for 
this choice will be given. If then the molar weight of a gas is 
proportional to its density the following relation will hold: 

The molar weight of the gas: the molar weight of oxygen = 
the density of the gas: the density of oxygen, or 

The molar weight of the gas The density of the gas 
The molar weight of oxygen The density of oxygen 

Substituting for the molar weight of oxygen its value 32, 
and solving for the molar weight of the gas, we have 

AT i n or> vx The density of the gas _ D 

Molar weight = 32 X T^T -3 rr ~r~ - = 32 X n nni/ton 

The density of oxygen 0.001429 

where 0.001429 is the density of oxygen and D that of the gas 
in question. 

But Q OQ142Q * s ^ e s P ec ifi g ray ity of the gas referred to oxygen. 
Since it is the ratio of the weights of equal volumes (1 c.c. in 
each case) of the gas and of oxygen. Therefore for a gaseous 
substance the molar weight may be defined as thirty-two times the 
specific gravity of the gas referred to oxygen as the standard. 

For the sake of illustration, let us calculate the molar weight 
of hydrogen. The density of hydrogen is 0.00008986, and 
letting M represent the molar weight, 

M = 32 X 0.00008986/0.001429 = 2.012 

Therefore the molar weight of hydrogen is 2.012. 

Water vapor of course cannot exist under standard conditions, 


but could exist at 0C. if the pressure were less than 0.4 cm. 
of mercury or at atmospheric pressure if the temperature were 
higher than 100C. Since water vapor follows the gas laws, 
if the weight of a given volume be found under any conditions of 
temperature and pressure, the density which it would have 
under standard conditions, if it did not liquefy, can be readily 
calculated. The density of water vapor obtained in this way 
is 0.0008063 and hence the 
Molar weight of water = 32 X - 000806 %.ooi429 = 18.05. 

By a slight extension of the conception of molar weight we 
can get another very useful conception, that of a gram mole. 
A gram mole of a substance is nothing more or less than the 
molar weight of the substance expressed in grams. Or it is a 
weight in grams of the substance numerically equal to its molar 
weight. For example, the molar weight of hydrogen is 2.012 
and a gram mole of hydrogen is 2.012 grm. of hydrogen. The 
molar weight of water vapor is 18.05, so a gram mole of water 
vapor is 18.05 grm. of water vapor. The molar weight of 
oxygen is 32, and a gram mole of oxygen then is 32 grm. of oxygen. 

The volume occupied by a gram mole of a gaseous substance 
under standard conditions can be easily calculated by dividing 
its molar weight by its density under standard conditions. The 
volume of a gram mole of oxygen under standard conditions is 
then 3 %.ooi43 or 22,400 c.c. The volume of a gram mole of 
hydrogen is 2 - 01 %.oooo8986' = 22,400 c.c. These and other 
examples which may be worked out, together with a few moments 
thought on the definition of molar weight should soon convince 
one that the volume of a gram mole of all gases under standard 
conditions is the same and is 22,400 c.c. 


There can be no question but that the law of combining weights 
is the most important generalization of chemistry, and we must 
now turn our attention to the formulation of this law. 
For this, we shall need the composition of a number of compounds 
as determined by analysis, and their molar weights. In order 
to obtain accurate results, it is necessary to use not the molar 
weights directly determined from the gaseous density, but these 
plus or minus a small experimentally determined number which 
corrects for the deviations of the gases from the simple gas laws 
of Boyle and Gay Lussac. These corrections are easily de- 
termined and the corrected molar weights are just as much the 
result of experiment as the uncorrected numbers. 1 

By analysis we get the weight of each element in a gram of the 
compound. If we multiply the weight of an element in a gram 
of compound by the molar weight of the latter, we obviously 
obtain the number of grams of .the element per gram mole of the 
compound. For example a gram of water contains according 
to analysis 0.8881 grm. of oxygen. The molar weight of water is 
18.016. If we multiply these two numbers together, we get 
16.00 which means that 18.016 grm. of water contains 16 grm. 
of oxygen. 

The following table summarizes this kind of results for a 
number of compounds, and should be carefully studied, paying 
particular attention to- the weights of each element per gram 
mole of compound. 

The results given in the following table under the headings, 
"weight of each element per gram and per gram mole of com- 
pound," are as described by the law of definite proportions and 

1 Some idea of the magnitude of these corrections may be obtained by 
comparing the value for the molar weight of water obtained on p. 69 with 
the corrected value given in the table p. 71 and also that uncorrected 
molar weight of hydrogen 2.012 with the corrected 2.016. 




each set represents equally well the relative proportions in which 
the elements unite to form the compounds. The most careful 
study of the weights of the elements per 1 gram of compound will 
fail to shows any simple relation between the weights of any given 
element, say hydrogen, in the various compounds; but even a 
very hasty inspection of the weights of the elements per gram 
mole shows that a gram mole of a hydrogen compound either con- 
tains 1.008 grm. or some integral multiple of this number. The 
same thing is true for oxygen, nitrogen, etc., except that the 
number is 16 for oxygen, 14.006 for nitrogen, 35.46 for chlorine, 
and 12 for carbon. 

Name of 

Weight of each ele- 
ment per gram of 


Weights of each element per gram mole of 








0.1119 hydrogen 
0.8881 oxygen 





0.0276 hydrogen 
0.9724 chlorine. 





0.1776 hydrogen 
0.8224 nitrogen 





0.4668 nitrogen 
0.5332 oxygen 





0.6365 nitrogen 
0.3635 oxygen 





0.2727 carbon 
0.7273 oxygen 





Minimum weight of 






each element per 

gram mole of the 


This is true not only for these compounds but also for all 
other gaseous compounds of these and other elements so that 
we can make the general statement that for each element there is 
a number which represents the minimum weight of that element 
found per gram mole of any of its compounds, and whenever this 
element is present in a gram mole of any compound in greater 
weight than the minimum, its weight may always be represented by 
an integral multiple of the minimum weight. These minimum 
numbers are called the combining weights of the elements because 
they or their integral multiples represent the proportions 
by weight in which the elements enter into combination with one 


another. The statement just given in italics might be called 
the law of combining weights, but if preferred the following 
form may be used. For each element a number can be found 
which has the property that the proportions by weight in 
which this element enters into combination with other elements 
can be represented by this number or some integral multiple 
of it. 

No two elements have the same combining weight and a 
table of these very important constants is given on p. 79. 

Atomic Theory. Just before the beginning of the nineteenth 
century, the law of definite proportions (p. 27) was formulated, 
and a few years later John Dalton, an Englishman, discovered 
some of the facts upon which the law of multiple proportions 
(p. 84) is based. From a consideration of these relations 
Dalton became aware of the facts described by the law of com- 
bining weights. He also knew the law of the conservation of 
the elements. As an explanation of all these facts he gave a 
definite quantitative form to the atomic theory a theory which 
had been held by many in some form or other since the days 
of the early Greek philosophers. For the purpose of explaining 
the facts, Dalton made the following assumptions: 

First, that all elements are made up of very minute chemically 
indivisible particles called atoms; 

Second, that the atoms of any one element are all alike in 
size, weight, etc., but are very different in all these particulars 
from the atoms of other elements; 

Third, that the ratio of the weights of the atoms of any two 
elements is the same as the ratio of their combining weights; 

Fourth, that compounds are formed by the union of a definite 
integral number of atoms of one element with a definite whole 
number of atoms of another to form the smallest particle, called 
the molecule, of the compound; 

Fifth, that the atoms of any given element preserve their 
identity in passing from the element to a compound, or from 
compound to compound. 

The explanation of the laws upon which the atomic theory is 
founded in terms of that theory is so simple as to be almost 
obvious. The fixed relative proportions by weight in which the 
elements unite to form a given compound (Law of Definite 


Proportions) is given by the relative weights of the atoms of 
the elements, each multiplied by the definite integral number of 
the atoms of that element which united to form a molecule 
of the compound. Since the atoms are assumed to be indi- 
visible, it follows that in a molecule of any given compound of an 
element, there will be one, two, three, or some whole number of 
atoms of that element; and therefore from the relation between 
the combining weights and the weights of the atoms, the propor- 
tions by weight in which the element enters into combination 
will be represented by the combining weight or some whole 
multiple of this (Law of Combining Weights). The law of 
conservation of elements is explained at once by the assumption 
that the elements preserve their identity in passing from com- 
pound to compound. In terms of the atomic theory, what we 
have called the combining weights of the elements may just 
as properly be called the relative atomic weights, or simply the 
atomic weights of the elements. The table on p. 79 entitled 
" International Atomic Weights" gives the accepted values of 
these important constants for the elements. 

The atomic theory has proved to be one of the most fruitful 
theories of chemistry, since it not only explains the very funda- 
mentally important laws of chemical combination, but also has 
brought together a great many other facts and by predicting 
new relations has led to investigations which have greatly 
increased our store of knowledge. For a long time there was no 
proof that substances were composed of particles the size of 
atoms and molecules; for all elements and chemical compounds 
when examined under the most powerful microscopes seemed to 
be perfectly homogeneous. But within the last few years, facts 
have come to light chiefly by the study of radioactive substances 
(see p. 496) which seem to indicate very clearly that such small 
particles do actually exist, and hence the theory has been greatly 
strengthened. These facts have even enabled us to make a 
very close estimate of the number of atoms in a gram combining 
weight or its synonym gram atomic weight of any element. 
This number is 6.1 X 10 23 atoms. 1 For example 1.008 grm. of 
hydrogen or 16 grm. of oxygen will each contain 6.1xl0 23 atoms. 

iG.l X 10 23 is an abbreviated way of writing 610,000,000,000,000,000,- 
000,000. - 


This almost inconceivably large number of atoms in a compara- 
tively small mass of substance enables us to understand fully 
why it has never been possible to see the atoms. 

Molecular Theory. Just as the atomic theory has proved very 
useful in explaining the facts of chemical combination, so the 
molecular theory has been of help in explaining certain of the 
other facts of chemistry and physics. Among the facts which 
the theory was devised to explain are those of the behavior 
of gases, liquids and solids, with changes of pressure and tem- 
perature, and also the relations which we have been considering 
under the conception of molar weights. 

The more important assumptions made for the purpose of 
explaining these facts are: 

First, that every substance is made up of extremely small 
particles called molecules. 

Second, that these molecules in the case of compounds are 
made up by the union of a definite and whole number of atoms 
of each of the -elements composing the compound ; in the case of 
elements, the molecules may be composed of one or more 
atoms usually more than one; 

Third, if anything happens to change the number of atoms in 
a molecule, the nature of the substance is changed; 

Fourth, that the molecules behave as though they were exceed- 
ingly elastic bodies in a state of rapid and ceaseless motion in 
all directions, their rates of motion increasing with rise of 

Fifth, because of their rapid motions, the molecules are 
assumed to occupy only a very small part of the space which is 
apparently filled by the substance which they constitute. The 
rest of the space being empty. This dispersion of the mole- 
cules is assumed to be especially great in the case of gases, much 
less for liquids, and generally still less for solids; 

Sixth, it is assumed (Avogadro's hypothesis) that equal 
volumes of all gases under the same conditions of temperature 
and pressure contain the same number of molecules. 

The pressure exerted by a gas is explained as being due to the 
hammering of the rapidly moving molecules upon the walls 
of the containing vessel. The latter part of the fourth assump- 
tion will explain the increase in pressure with rising temperature, 


if the volume is kept constant; or the increase in volume, if the 
pressure is unchanged. (Gay Lussac, or Barton's law.) 

Decreasing the volume occupied by a given mass of a gas to 
one-half its former value will double the number of molecules 
striking a given area of the wall confining the gas and hence 
double the pressure. (Boyle's law.) 

If Avogadro's hypothesis is true and equal volumes of all gases 
under the same conditions of temperature and pressure, contain 
the same number of molecules, then it follows that the weights 
of these equal volumes will stand to each other in the same ratio 
as the actual weights of the molecules. We can then by choosing 
some gas as the standard and assigning to it a certain molecular 
weight, prepare a system of relative molecular weights of gases 
which shall have all the properties of that which we have called 
the molar weights. In fact since oxygen has been taken as the 
standard and assigned a molecular weight of 32, the numerical 
values of the molecular weights is just the same as that for molar 
weights; and a gram mole of a substance is precisely the same 
quantity of that substance as a gram molecular weight. The 
expression, molecular weight, is a rather long one, so in accord- 
ance with custom, it will generally be replaced in this book by 
the more convenient word mole. 

The same facts of radioactivity (p. 496) which strengthened 
the atomic theory support the molecular theory and indicate 
that the number of molecules in a gram molecular weight or 
gram mole of a substance is the same as the number of atoms 
in a gram atomic weight that is to say 6.1 X 10 23 . A further 
strengthening of the molecular theory is furnished by the fact that 
extremely minute particles, when suspended in a liquid and 
strongly illuminated, look, when examined by a powerful micro- 
scope, like dancing dots of light. The motions of these particles 
closely correspond in kind and amount to those calculated on the 
assumption that they are caused by the bombardment of the 
particles by bodies having the calculated mass and velocity of 
the assumed molecules. 

Relation Between Fact and Theory. In studying a theory we 
should always be careful to get thoroughly in mind the distinc- 
tion between the facts, the assumptions and the explanation. 
The facts are eternally true. The theory is a speculative ex- 


planation of the facts in which our imagination has carried us 
farther than our observations have been able to go. Further 
observations may show that our imaginings (the theory) was a true 
picture of the way nature works, or they may bring to light new 
facts which are so out of harmony with what we had imagined 
that we have to abandon our former theory and invent an en- 
tirely new explanation which shall at the same time cover the 
old and the new found facts. This last condition has been the 
usual one in the past and the history of science shows a great pile 
of discarded theories. Each of these was a good theory in its day 
because it explained the known facts and thus simplified and 
coordinated our knowledge and also predicted the existence of 
other facts and so stimulated investigation with the result that 
our field of information was greatly extended. Judging from 
the past we have every reason to expect that the future will 
replace many of our present theories by better ones. The only 
difficulty about a false theory is that if it is too firmly believed 
in, it blinds one to obvious facts and causes others to be mis- 
understood, and so hinders the progress of science. 

The way to use a theory is to keep constantly in mind that 
it is not a fact and possibly is not true, but to follow up all the 
lines of investigation which it suggests, always being ready to 
discover that the facts do not agree with the theory. In particular, 
one should never hesitate to try a thing which upon other 
grounds promises success simply because a theory predicts 

Standard of Atomic Weights. Our recently acquired knowl- 
edge of the number of atoms in a gram atomic weight would 
enable us to prepare and use a table of the absolute values of the 
atomic weights instead of that given on p. 79, but these numbers 
would be long decimal fractions with from 21 to 22 ciphers before 
the significant figures, and hence would be far less convenient than 
our present system which like that for molar weight is entirely 
relative, and therefore necessitates the arbitrary choice of a 
standard. At first, hydrogen was taken as the standard because 
it has the smallest combining weight, and was given the relative 
weight of 1 ; but of recent years it has been decided that oxygen 
with a relative weight of 16 is the better standard. One of the 
principal reasons for this change to oxygen is that many elements 


do not form easily volatile compounds so that it is very difficult 
to determine their combining weights from their minimum weight 
per gram mole of gaseous compounds. ' Under these conditions, 
it is necessary to determine the proportions by weight in which 
the element of unknown atomic weight will combine with one 
for which this constant is well known preferably the standard. 
Now hydrogen does not form stable compounds with many of 
the other elements, while oxygen does. Therefore oxygen is the 
better for a standard substance. Of course the weight of an 
element which will combine with 16 grm. of oxygen will not give 
us directly the atomic weight of the element unless just one atom 
of it combines with one atom of oxygen; but in any event the 
weight per 16 grm. of oxygen bears some very simple relation to 
the atomic weight i.e., it is twice the atomic weight if two atoms 
of the unknown combine with one of oxygen, or one-half it if two 
atoms of oxygen combine with one atom of the other. There is 
a rather inexact law known as that of Dulong and Petit which 
states that the product of the atomic weight of a solid element 
and its specific heat expressed in calories is approximately equal 
to the constant 6.2. The deviations from this number are rather 
large, especially for the non-metallic elements of low atomic 
weight, but in nearly all cases, the value of the atomic weight 
obtained by determining the specific heat and dividing this into 
6.2 is nearly enough correct, so that it will enable us to determine 
the real value, from the weight which will combine with 16 grm. 
of oxygen. For example, the specific heat of magnesium is 0.245 
and this divided into 6.2 gives 25.3 as the approximate atomic 
weight of magnesium. But 24.32 grm. of magnesium will com- 
bine with 16 grm. of oxygen and therefore the atomic weight of 
magnesium must be 24.32. For silver the case is a little more 
complex. The specific heat of silver is 0.057 and this divided 
into 6.2 gives 109 as the approximate atomic weight of silver. 
But 215.76 grm. of silver will combine with 16 grm. of oxygen. 
Since 215.76 is so nearly twice 109, it is evident that in this case 
two atoms of silver combine with one of oxygen, and that the 
atomic weight of silver is 215.7% = 107.88. It cannot be too 
strongly emphasized that these atomic weights are ratio numbers, 
and that they or their integral multiples represent the propor- 
tions by weight in which the elements combine. 


Symbols and Formulas. In order to save time and also for 
the sake of securing a convenient method for the representation 
of the composition of substances, a system of symbols has 
been devised which shall at the same time represent the elements 
and the atomic weights of the elements, and which by com- 
bining the symbols of the elements shall represent the compounds, 
the molecular weights of the compounds and their composition. 
This system was devised by the celebrated Swedish chemist 
Berzelius. The symbols for the elements are derived from 
the names of the elements, using the Latin or Greek word where 
there is a difference in the name as we pass from language to 
language. In most cases the symbol is the first letter of the 
name, for example, the symbols of oxygen and hydrogen are 
O and H respectively. In case the names of more than one 
element begin with the same letter, the symbol of the element 
is formed from this initial letter and a characteristic letter from 
the name. Take for example chlorine with the symbol Cl. 
The first letter of the symbol is always a capital and the second 
is always a small letter. In addition to representing the element 
one of these symbols also stands for an atomic weight of the ele- 
ment, and so represents a certain number of parts by weight of 
the element. O stands for oxygen, and 16 parts by weight of 
oxygen; H for hydrogen, and 1.008 parts by weight of hydrogen; 
Cl stands for chlorine, and 35.46 parts by weight of chlorine. 

Hydrogen and chlorine combine to form hydrogen chloride 
in the proportion of 1.008 parts by weight of hydrogen to 
35.46 parts by weight of chlorine, or of one atomic weight of 
the one element to one of the other. This is the composition 
of hydrogen chloride and both hydrogen chloride and its com- 
position can be represented by writing the symbols of the ele- 
ments close together, viz., HC1. This HC1 stands first for 
hydrogen chloride; second for a mole of hydrogen chloride, that 
is, 36.468 parts by weight of the substance; third it shows that 
a mole of hydrogen chloride is formed by the combination of 
one atomic weight of each of the elements, or of 1.008 parts by 
weight of hydrogen to 35.46 parts of chlorine. The sum of these 
atomic weights gives the molecular weight or the parts by weight 
of the hydrogen chloride represented by the formula HC1. 

Water is formed by the union of 2.016 parts by weight of 















20 2 






58 68 



137 37 

Niton (radium ema- 


222 4 






14 01 










Oxygen . 


16 00 

Calcium . . 





106 7 

Carbon. . 


12 005 



31 04 




Platinum. ... ... 








39 10 















93 1 



85 45 



63 57 



101 7 







Erbium. . 



Scandium . 








79 2 






28 3 








. Ga 

69 9 

Sodium . . . 





72 5 



87 63 



9 1 



32 06 



197 2 






4 00 

Tellurium ... . 





1 008 




Indium. . 



Thallium '. 





126 92 



232 .4 



193 1 






55 84 

Tin. . 





82 92 

































200 6 










hydrogen with 16.000 of oxygen to form 18.016 parts by weight 
of water. The 2.016 parts by weight of hydrogen is two atomic 
weights of this element, while the 16 is of course only one atomic 
weight of oxygen. To represent the composition of water, <then, 
we must build up a formula which shall stand for two atomic 
weights of hydrogen and one of oxygen. We might write it 
HHO or HOH, and it is sometimes written the latter way. The 
almost universal method of writing it though is H 2 O. The sub- 
script figure 2 being used to show the number of atomic weights 
of hydrogen per mole, the or any symbol used without the 
subscript being understood to represent one atomic weight of 
the element. The formula H 2 O then stands for water, 18.016 
parts by weight of water, or a mole of water, and shows that two 
atomic weights of hydrogen unite with one atomic weight of 
oxygen to form a mole of water. 

Great care must be taken by the student to 'avoid getting 
the notion that the subscripts represent the proportions by 
volume in which the elements combine. It so happens in the 
instances cited above that this is the case, but it is by no means 
always true, and one should never permit himself to think of 
it in this way. 

As an exercise the student should write the formulas for the 
rest of the compounds whose compositions are in the table on 
page 71, and also write out in words just what each represents. 

The preceding table gives the latest values for the combining 
or atomic weights. 

Molecular Formulas of the Elements. The following table 
gives the atomic and molecular weights, together with the 
formulas of a number of typical elements. 

Name of element 




Number of 
atoms per 




O 2 



1 008 

2 016 

H 2 






N 2 





C1 2 


Mercury. . ... 






31 04 


P 4 



The formulas of the elements are obtained by applying the 
principles in the section on Symbols and Formulas. An inspec- 
tion of the third and fourth columns of the table shows that the 
atoms of an element are usually not identical with its molecules, 
but that one of the latter is generally made up by the union of 
two or more atoms. For example, the molecule of oxygen con- 
tains two atoms of oxygen and has a formula of 02. At first 
sight, it might appear that this is due to our choice of oxygen 
as the standard substance for both atomic and molecular 
weights, and arbitrarily giving it an atomic weight of 16 and a 
molecular weight of 32; and that if we had 'called it 16 in each 
case, that the atom and the molecule would have been identical. 
But the following reasoning will show that the atom cannot be 
the same as the molecule. Hydrogen and oxygen combine to 
form water in the proportion of two volumes of hydrogen to one 
of oxygen, (see p. 67) forming two volumes of water vapor. 
If we take as the unit of volume the volume of one gram mole 
22,400 c.c. under standard conditions we will have the following 
volume and weight relations: 
2 X 22,400 c.c. of hydrogen + 22,400 c.c. oxygen = 2 X 22,400 

c.c. water vapor. 

2 X 2.016 grm. of hydrogen + 32 grm. oxygen = 2 X 18.016 
grm. water vapor. 

Then 2 grm. moles of hydrogen react with 1 grm. mole of oxygen 
to form 2 gram moles of water vapor; but each gram mole of a 
substance on the basis of Avogadro's hypothesis contains the 
same number of molecules. Therefore there are just twice as 
many molecules of water as there are molecules of oxygen. 
The assumed indivisibility of atoms makes it impossible to assume 
that a molecule of water contains less than one atom of oxygen; 
hence if a molecule of water contains one atom of oxygen, a 
molecule of oxygen must contain two atoms. Therefore if the 
atomic weight of oxygen is taken as 16, its molecular weight must 
be 32. This then is the basis for the apparently arbitrary choice 
of 32 as the molar weight of oxygen (p. 68). 

Chemical Equations. As was said above, two atomic weights 
of hydrogen unite with one of oxygen to form a mole of water; 
this can be very well shown with the aid of symbols as follows; 

H 2 + O = H 2 O 


and this is an example of a chemical equation. A chemical 
equation is a very simple shorthand method of representing 
a chemical reaction quantitatively. For example, the above 
equation shows that 2.016 parts by weight of hydrogen will 
combine with 16 parts by weight of oxygen to form 18.016 
parts by weight of water. The unit of weight may be the gram, 
pound or any other weight unit, and the equation will hold 
just the same. 

We have learned that two volumes of hydrogen will combine 
with one volume of oxygen to form two volumes of water vapor, 
or two moles of hydrogen and one mole of oxygen will give two 
moles of water vapor. In representing this we shall have to use 
the molecular formulas for everything. Now the molecular 
weights of hydrogen and of oxygen are twice their atomic weights' 
so their molecular formulas are H 2 and O 2 respectively. We 
may show that two moles of hydrogen combine with one of oxygen 

2H 2 + O 2 = 2H 2 O 

This is read "two moles of hydrogen combine with one mole 
of oxygen to give two moles of water vapor." The coefficients 
or numbers before the formulas show the number of moles taking 
part in the reaction, while the subscripts show the number of 
atomic weights of that element per mole. The coefficients 
therefore multiply every symbol in the formula which they 
stand before, as for example, 2H 2 represents 2 moles of water 
of 2 X 2 atomic weights of hydrogen and 2X1 atomic weights of 
oxygen. On either side of the above equation there are repre- 
sented then, four atomic weights of hydrogen and two of oxygen. 
This is the proof that the equation is properly "balanced." 
There are two main principles involved in writing these equa- 
tions. First, the formulas separated by + should be molecular 
formulas, and second, there must be the same number of atomic 
weights of any given element on either side of the equation, since 
otherwise we would be indicating a deviation from the law of 
the conservation of weight. The advantage of writing molecular 
formulas lies in the fact that since gram moles of gases occupy 
equal volumes under like conditions we can tell at once their 
proportion by volume, from a simple inspection of the number 


of moles of the one that reacts with the other. The proportions 
by volume in which gases react are given by the coefficients in 
the molecular equation for the reaction. 

When one of these equations is first worked out, it must be 
found by experiment just what things are formed by the sub- 
stances which react and the proportions by weight of everything 
involved in the process. There is no other way of doing it. We 
must know what is brought together and what is formed before 
we can represent it by symbols. Students often think that they 
fail to grasp the idea of equations because they cannot write 
down symbols and formulas for a number of substances on one 
side of the equation and then tell by some rule what to put on 
the other, or else they jump to the other extreme and think that 
the symbols can be combined in any way at the fancy of the 
writer. But the truth of the matter is that only certain sub- 
stances are formed from certain other ones, and one must know 
what these are and then represent them. Every equation then 
represents the results of actual experiment. It will be recalled 
that water and sodium react and form sodium hydroxide and 
hydrogen. The sodium hydroxide contains per mole 23 parts by 
weight of sodium, 16 of oxygen, and 1.008 of hydrogen. Its 
formula then is NaOH and the molecular weight is a little over 
40. We find by experiment that when 36.032 grm. of water 
react on 46 grm. of sodium that 2.016 grm. of hydrogen and 
80.016 grm. of sodium hydroxide are formed. We may represent 
this by the following equation, 

2H 2 O + 2Na = H 2 + 2NaOH 

In studying the properties of substances it is very essential to 
learn with what other substances they will react and the sub- 
stances formed. It is not necessary to learn in what propor- 
tions they react as, given the other information, the proportions 
can be reasoned out from the rules for balancing which were 
given above. 

Knowing the substances and their weights we can write the 
equation, or given the equation we can work out the weights 
from the atomic weights. 


Hydrogen and oxygen will combine in the ratio of one part by 
weight of hydrogen to 7.943 parts by weight of oxygen, and in 
no other ratio when they unite to form water. However, they 
will combine in a different proportion for the formation of an 
entirely distinct substance known as hydrogen peroxide. The 
ratio in this case is one part by weight of hydrogen to 15.886 parts 
by weight of oxygen. If we now compare the parts by weight 
of the oxygen in the two compounds which will combine with 
one part of hydrogen we see that they stand to each other as 1 : 2 

7.943 : 15.886 = 1:2 

This is the first time that we have found a really simple weight 
relationship and this you will notice is the ratio between the 
weights of the one element which will combine with unit weight 
of the other element in the two compounds. This sort of a simple 
weight relation has been found to be so common that a law has 
been formulated to describe it. 

The Law of Multiple Proportions. // two elements, which we 
may represent by A and B, combine to form more than one com- 
pound then the number of grams of A which combine with one 
gram of B in the first compound will stand to the number of grams 
of A per gram of B in each of the other compounds in the propor- 
tions of whole numbers and usually these numbers are small ones. 

Hydrogen and oxygen always combine in one fixed ratio for 
the formation of water, and in another equally fixed for the for- 
mation of hydrogen peroxide, so the law of definite proportions 
describes the composition of each of these compounds and the 
law of multiple proportions applies to the two taken together. 

Preparation and Properties of Hydrogen Peroxide. Hydrogen 
peroxide is prepared in a number of ways, but the greater part 
of that used is made by the action of dilute sulfuric acid upon 
hydrated barium peroxide, barium sulfate being formed at the 



same time. It is not practical to prepare it directly from the 
elements although minute quantities are formed by allowing the 
flame of an oxy-hydrogen blowpipe to play upon ice. The 
hydrogen peroxide prepared by the interaction of barium 
peroxide and sulfuric acid is present in a dilute solution with 
water. To prepare pure hydrogen peroxide is not easy, but the 
method usually adopted depends upon the fact that hydrogen 
peroxide is less volatile than water, so that if a mixture of the 
two be distilled, relatively more of the water than of the hydro- 
gen peroxide will pass off, and the residue will contain more and 
more hydrogen peroxide per cubic centimeter. By repeating the 
process, nearly pure hydrogen peroxide may be obtained. 
Hydrogen peroxide is rather unstable and decomposes into 
water and oxygen. This reaction like nearly all others goes 
more rapidly at high temperatures than at lower ones, there- 
fore the process of distillation is carried out under diminished 
pressure because this lowers the boiling-point. 

Hydrogen peroxide is a rather thick, syrupy liquid which has 
a very faint greenish-blue color somewhat more intense than that 
of water. Its density at 0C. is 1.46. It melts at -2 and 
boils at 69 under a pressure of 2.6 cm. of mercury. Its boiling- 
point under atmospheric pressure is not known as it decomposes 
with explosion before it boils. It is odorless and has an astrin- 
gent bitter taste. It is soluble in water in all proportions and 
the solution looks like water. This solution slowly decomposes 
into water and oxygen and keeps best in the dark, in a cool place, 
and when it contains a little acid. This transformation seems 
especially sensitive to the presence of foreign substances which 
after the completion of the process are found unchanged. These 
substances are therefore catalyzers. The following is a partial 
list of the catalyzers of the decomposition of hydrogen peroxide ; 
charcoal, silver, gold, platinum, manganese dioxide, potassium 
and sodium hydroxides, saliva, and blood. The solid substances 
mentioned act the more vigorously the finer they are powdered. 

A 3 per cent, solution is a standard article of commerce. It 
is often called a 10-volume solution from the fact that it will 
evolve, when all the hydrogen peroxide is decomposed, about 
ten times its volume of oxygen. This solution will rapidly 
destroy bacteria, and hence is much used as a disinfectant. 


Oxidizing Action of Hydrogen Peroxide. Since hydrogen per- 
oxide gives up part of its oxygen so easily that it will do it even 
at ordinary temperatures, it seems very natural to expect that 
it will be able to give it up to other things which can combine with 
oxygen and thus oxidize them. Upon trial this is found to be 
the case, and this action is exhibited more strongly as the 
hydrogen peroxide becomes more concentrated; that is, the 
greater the weight of this substance per cubic centimeter. If the 
pure hydrogen peroxide be brought in contact with many in- 
flammable substances, the 'latter are oxidized so rapidly that 
they ignite spontaneously. The solutions are much less intense 
in their action and decrease in activity as the dilution is 

Substances which will give up oxygen to other substances 
in the way that hydrogen peroxide does are called oxidizing 
agents. As we shall see later, oxidizing agents can do many 
things besides give up oxygen to other substances, but this is 
one of the things which they may do, and must be kept in mind 
until finally we come to a definition of an oxidizing agent which 
will cover all cases. Other oxidizing agents are like hydrogen 
peroxide, in that they act more vigorously as their concentration 
increases. Because of its action as an oxidizing agent, hydrogen 
peroxide is able to change many organic coloring matters into 
colorless substances, and it is therefore used for bleaching hair, 
feathers, ivory, silk, wool, bones, leather, etc. For this purpose 
fairly dilute solutions are used. 

Reducing Action of Hydrogen Peroxide. When hydrogen 
peroxide acts upon many oxides such as those of silver, gold 
and mercury, the metals, water and oxygen are produced. The 
process of obtaining a metal from its oxide is an example of 
reduction, and hence hydrogen peroxide acts in these cases as 
a reducing agent. Potassium permanganate is also reduced by 
hydrogen peroxide and in this case again oxygen is evolved. 

Determination of the Molecular Weight of Hydrogen Peroxide. 
Hydrogen peroxide is so unstable that it is not possible to 
determine its molecular weight from its gaseous density, but re- 
course must be had to an entirely different method called the 
' ' Freezing-point Method . ' ' 

It will be recalled that mention has been made of the fact 


that any dissolved substance lowers the freezing-point of water, 
and also that gram moles of the various substances produce equal 
physical effects. Experiment has shown that if a gram mole 
of substance be dissolved in 1,000 grm. of water, the freezing- 
point of the water is lowered 1.85C. This then furnishes a 
ready method for the determination of the molecular weight of 
such substances as hydrogen peroxide which cannot be vapor- 
ized without decomposition. The molecular weight of a sub- 
stance, with some exceptions, as noted below, is identical with 
the number of grams of the substance required to lower the 
freezing-point of 1,000 grm. of water 1.85C. 

Any convenient weight of the solute and of water may be 
used and the molecular weight of the solute calculated from 
the equation, 

M = 1,850^ 

in which "M" represents the molecular weight of the solute, 
"w" the weight of the solute, "W" the weight of the water, 
"L" the lowering of the freezing-point. 

The molecular weight of hydrogen peroxide determined in this 
way comes out close to 34. 

Under the discussion of water it was pointed out that solutes 
lowered the vapor pressure of their solvents and that there was 
a connection between this fact and the fact that the solutions 
froze at lower and boiled at higher temperature than their 
pure solvents. From these facts, one would be led to expect 
that there would be a constant rise in the boiling-point of water 
per gram mole of solute per 1,000 grm. of water, because gram 
moles are in the main, physically equivalent. This is found to be 
the case. One gram mole of a substance with the exceptions 
noted below will raise the boiling-point of 1,000 grm. of water 
0.52C., and a method for the determination of molecular 
weights is based upon this. The formula is 


M = 


in which M, w, and W have the same significance as before, and 
"r" is the rise in the boiling-point, and 520 the proportionality 


As has been implied above, there are some exceptions to the 
rule that a gram mole will lower the freezing-point of 1,000 grm, 
of water 1.85C or raise the boiling point 0.52C. We will 
find later that these exceptions lead to very interesting and im- 
portant conclusions. All these exceptional substances yield 
solutions which are conductors of electricity. Hydrogen 
peroxide in solution is a non-conductor, so we may feel con- 
fident that the molecular weight given above is correct. 

Composition and Formula of Hydrogen Peroxide. Before 
the formula of any compound can be worked out the substance 
must be analyzed or synthesized and its molar weight determined. 
From the data the formula is then made up in a way which 
may be well illustrated by hydrogen peroxide. 

The analysis of hydrogen peroxide shows that it consists 
of 5.93 per cent, hydrogen and 94.07 per cent, oxygen. 

From these data and its molecular weight 34.016, it is a simple 
matter to calculate the number of parts by weight of each 
element per mole of the peroxide as follows: 

34.016 X 5.93 

- = 2.016 parts by weight of hydrogen 

34.016 X 94.07 00 n . , ^ f 

- = 62.0 parts by weight 01 oxygen. 


Since the atomic weight of hydrogen is 1.008 and of oxygen is 
16, it is evident at a glance that there is in a mole of hydrogen 
peroxide two atomic weights of hydrogen and two of oxygen 
and therefore its formula is H 2 2 . 

Equation for the Preparation of Hydrogen Peroxide. As has 
been already mentioned, hydrogen peroxide may be prepared by 
the action of dilute sulfuric acid upon barium peroxide, barium 
sulfate being formed at the same time. The equation for the 
reaction is as follows: 

BaO 2 + H 2 SO 4 = BaSO 4 + H 2 O 2 

in which Ba0 2 stands for 169.37 parts by weight of barium 
peroxide, H 2 S0 4 for 98.076 parts by weight of sulfuric acid, BaS0 4 
for 233.43 parts by weight of barium sulfate, and H 2 2 for 34.016 
parts by weight of hydrogen peroxide. 


Explosive Properties of Hydrogen Peroxide. Pure or nearly 
pure hydrogen peroxide is decidedly explosive. The expla- 
nation for this is that the peroxide passes into water and oxygen 
with the evolution of a great amount of heat, so that if decompo- 
sition gets started, the heat evolved raises the temperature of 
the remaining peroxide and this makes it break up more rapidly, 
evolving still more heat and raising the temperature yet higher, 
thus making it decompose even faster until the reaction takes 
place so violently that it is explosive. All other explosives are 
like hydrogen peroxide in that they give off heat during their 
transformation, and in fact they owe their explosiveness to this 

Explosives are rather unusual substances because most things 
decompose with the absorption of heat, and hence if they once 
start to decompose they cool themselves off and so tend to stop 
the action. So pronounced is the heat of decomposition of 
hydrogen peroxide that even the dilute solution of commerce will 
experience a rise in temperature of 15 to 20C. upon the addition 
of manganese dioxide which will catalytically bring about a rapid 
destruction of the peroxide. 

Thermochemistry. Practically every chemical change is ac- 
companied by either the absorption or evolution of heat. These 
heat effects are of importance because, if the conditions are 
made such that no work is done, they represent the difference 
in energy between the substances before and after the reaction. 
These differences in energy are of vital importance to us because 
it is really the difference in energy between the coal and oxygen 
on the one hand and the products of combustion on the other 
which drives the steam engine, cooks our food, and keeps us 

These heat effects are measured in very much the same way 
as the heat of fusion of ice or the heat of evaporation of water, 
and are expressed in calories or kilo joules per gram, or better 
still per gram mole of substance transformed. 

By a very slight addition to our scheme for representing 
chemical reactions we can indicate at the same time the chemical 
and energy changes. In doing this the chemical formulas stand 
for gram moles of the various substances, and at the end of the 
equation is written the number of kilo joules "Kj" or calories 


evolved or absorbed, evolution being indicated by a plus (+) 
sign and absorption by a minus ( ). 

To illustrate this, let us consider the heat effect accompanying 
the burning of hydrogen to form water. For every gram mole 
of water produced 68,400 calories or 286 kilojoules of heat are 
evolved. This is represented as follows: 

2H 2 + 2 ~ 2H 2 + 2 X 68,400 cal. 
or 2H 2 + O 2 = 2H 2 + 2 X 286 Kj. 

The heat effect accompanying the formation of a gram mole of 
a substance is called its heat of formation, so the heat of forma- 
tion of water is 286 Kj. or 68,400 cal. per mole. 

There are two points in connection with thermochemical 
equations such as those just given, which must be carefully kept 
in mind. The first one is that if a certain quantity of heat is 
evolved during the formation of a gram mole of a substance, pre- 
cisely the same quantity of heat will be taken up during its 
decomposition. The second is that the heat effect accompanying 
a given transformation depends only upon the initial and final 
states of the substances and not at all upon the intermediate 
stages through which they pass. 

The Heat of Formation of Hydrogen Peroxide. Since hydrogen 
and oxygen combine with such difficulty to produce hydrogen 
peroxide it is not possible to determine the heat of formation 
of this substance by direct measurement. This magnitude may, 
however, be calculated by taking advantage of the facts pointed 
out above. Hydrogen peroxide spontaneously decomposes into 
water and oxygen, and the heat effect belonging with this change 
may be easily measured. It is 97 Kj. or 23,100 cal. per mole of 
peroxide transformed. The equation is 

2H 2 2 = 2H 2 + 2 + 2 X 97 Kj. (I) 

From what has been said above and the law of the conservation 
of energy, it follows that if oxygen and water should unite to form 
hydrogen peroxide that 97 Kj. of heat would be absorbed for 
each gram mole of peroxide formed, the equation being 

2H 2 + 2 = 2H 2 2 - 2 X 97 Kj. (II) 

The heat of formation of the peroxide from oxygen and water 


therefore is 97 Kj. To obtain the heat of formation from the 
elements, all that is necessary is to add to the above number 
the heat of formation of the water. In fact these thermochemical 
equations may be handled like ordinary algebraical equations. 
To equation (II) may be added the equation for the formation 
of ( water and the result will be the equation for the formation 
of the peroxide from its elements. 

2H 2 + 2 = 2H 2 2 - 194 Kj. (II) 

2H 2 + 2 = 2H 2 O + 572 Kj. 
2H 2 + 2O 2 + 2H 2 = 2H 2 O 2 + 2H 2 + 378 Kj. 
or 2H 2 + 20 2 = 2H 2 2 + 2 X 189 Kj. 

So the heat of formation of hydrogen peroxide is 189 Kj. or 
45,300 cal. per gram mole. Of course the thermochemical 
equations expressed in cal. may be substituted for those in Kj. 
and the figure 45,300 will then be obtained directly. 

The method applied here is a general one for calculating the 
heat effects which cannot be directly measured. 


The third element which we will discuss is chlorine. This 
element does not occur free in nature but is found in very large 
quantities in compounds, such as sodium, potassium, and mag- 
nesium chlorides. The element itself and many of its compounds 
are of very great importance not only in the chemical industries, 
but also in every-day life. The sodium chloride, NaCl for 
example, is our common salt, and we all know how indispen- 
sable this is. Sodium chloride is found in enormous quantities 
in beds of rock salt at many places on the earth's surface and 
is present in still larger quantities in the ocean. 

Chlorine was first prepared by Scheele in 1774, but that it was 
an element was first shown by Sir Humphrey Davy in 1810. 

Preparation of the Element. There are a few chlorides, such 
as those of gold and platinum, which will decompose into chlorine 
and the metals upon being raised to a high temperature, but these 
chlorides are so expensive and difficult to make that chlorine is 
practically never prepared in this way. The preparation of 
chlorine, 'then, resembles that of hydrogen more nearly than that 
of oxygen, in that a compound of chlorine is treated with some 
substance which will convert the rest of the chlorine compound 
into a substance which will readily take on a different physical 
state from that of the gaseous chlorine and hence may be 
easily separated from this element. Chlorine may also be liber- 
ated by electrolysis of solutions of soluble chlorides ver^y much 
as hydrogen and oxygen may be prepared. In the case of the 
electrolysis of the chlorides, chlorine is given off at one electrode, 
and either hydrogen or the metal of the chloride is deposited at 
the other. 

The most suitable substance for the chemical preparation of 
chlorine seems to be hydrochloric acid, a compound of chlorine 
and hydrogen whose formula is HC1. From this hydrogen 
chloride or hydrochloric acid, we can easily prepare chlorine by 




converting the hydrogen into water. This may be done in several 
ways. Gaseous hydrogen chloride mixeol with oxygen or air may 
be passed through a tube containing clay balls or pieces of pumice 
stone coated with copper chloride or sulfate, heated to 370 to 
400C. Under these conditions water and chlorine are produced 
according to the equation, 

4HC1 + 2 = 2H 2 O + 2C1 2 

The copper salt acts as a catalyzer. 

When air is used the product is of course greatly contaminated 
with nitrogen, as well as unchanged hydrogen chloride and 
oxygen, so that the method is not suitable for laboratory purposes, 
but was formerly largely used for the commercial preparation of 
chlorine, and is known as the Deacon Process. 

Other Methods. The preparation of chlorine from hydrogen 
chloride belongs to the processes which we call oxidation, and 

FIG. 19. 

oxygen here is the oxidizing agent. Other oxidizing agents, if 
sufficiently powerful, may be used in the place of oxygen. Among 
those which are convenient may be mentioned manganese dioxide, 
Mn0 2 , potassium chlorate* KC1O 3 , potassium permanganate, 
KMn0 4 , and bleaching powder, CaCl 2 0. Each of these sub- 
stances when treated with hydrochloric acid will oxidize the 
latter to water and chlorine. 


A very commonly used laboratory method for the preparation 
of chlorine is to gently heat a mixture of manganese dioxide 
with four times its weight of commercial hydrochloric acid (Fig. 
19). The liquid first turns brown and then gives off greenish- 
yellow gaseous chlorine. After the reaction is finished, the 
solution contains manganous chloride, MnCl 2 . The equation is 

MnO 2 + 4HC1 = CU + MnCl 2 + 2H 2 

It seems to be very probable that the reaction goes in two stages, 
first forming water and manganese tetrachloride, MnCl 4 , which 
is dark colored, and then this tetrachloride decomposes into 
the lighter colored manganous chloride, MnCh, and chlorine, C1 2 . 
The equations representing this would be 

MnO 2 + 4HC1 = MnCU + 2H 2 
and MnCU = MnCU + CU 

Instead of starting with hydrochloric acid and manganese 
dioxide one may form the hydrochloric acid from salt and sul- 
furic acid and oxidize it with manganese dioxide right in the 
same vessel. To do this one warms a mixture of common salt 
and manganese dioxide with moderately dilute sulfuric acid. 
The chlorine comes off very regularly, and at a convenient rate, 
so that this modification of the first method is often used. The 
equation is 

2NaCl + MnO 2 + 3H 2 SO 4 = MnSO 4 + 2NaHSO 4 + C1 2 + 2H 2 O 

Perhaps the most convenient laboratory method for the prepa- 
ration of chlorine is to allow hydrochloric acid to fall drop by 
drop upon solid potassium permanganate, KMn0 4 . The action 
takes place promptly and at ordinary temperatures. Water, 
manganous chloride, potassium chloride, and chlorine are the 
products. The equation is 

2KMn0 4 + 16HC1 = 8H 2 O + 2MnCl 2 + 2KC1 + 5CU 

A very important method which is much used both in the labora- 
tory and technically is to treat bleaching powder, CaCl 2 0, 
with hydrochloric acid. Water, calcium chloride, CaCl2, and 
chlorine are the products. The equation is 

CaCl 2 + 2HC1 = H 2 + CaCl 2 + CU 


It should be noted that in each of the above cases liquid water, 
a non-volatile salt or salts, and gaseous chlorine are formed. This 
difference of state makes it easy to isolate the chlorine. 

Technical Preparation. Each of the above-mentioned methods 
for the preparation of chlorine, together with many others, has 
been used on a manufacturing scale, but at the present time 
the greater part of the chlorine of commerce is produced by the 
electrolysis of solutions of sodium or potassium chloride see 
Fig. 52, p. 281. The anode or positive electrode where the 
chlorine appears is usually made of graphite because almost 
everything else is attacked by the chlorine. At the cathode or 
negative electrode hydrogen is given off and sodium or potassium 
hydroxide is formed in solution. The process takes place as 
though metallic sodium or potassium were liberated at the 
cathode and then at once reacted with the water for the formation 
of hydrogen and sodium or potassium hydroxide. 

Physical Properties. At ordinary temperatures chlorine is 
a greenish-yellow gaseous substance whose density is 0.0032215 
and whose uncorrected molecular weight is 72.13. Chlorine 
does not follow the gas laws very closely, a fact which is connected 
with the circumstance that it is very easily converted into a liquid 
either by cooling it to 33.6C. under a pressure of one atmos- 
phere, or by putting it under a pressure of 3.66 atmospheres at 
0C. or of 6.62 atmospheres at 20C. The critical temperature 
is 146 and the critical pressure 94 atmospheres. After correcting 
for its deviations from the gas laws the molecular weight of 
chlorine is found to be 70.92. 

Liquid chlorine has a greenish-yellow color, a density of 1.56, 
and is of an oily consistency. At - 102 the liquid freezes to 
yellow crystals of solid chlorine. 

Gaseous chlorine is soluble in about half its volume of cold 
water but is much less soluble in hot. Because of its solubility 
in cold water and the fact that it attacks mercury with great 
vigor, chlorine must be collected over hot water or by displace- 
ment of air. The latter is the more convenient, and since chlorine 
is nearly 2.5 times as heavy as air, it is easily carried out. 

Chlorine has a very disagreeable odor, is exceedingly irritating 
to the mucous membrane of the throat and nose and has often 
produced death. 


Chemical Properties. The atomic weight of chlorine is 35.46 
and its corrected molecular weight is twice this, so there are 
two atomic weights per mole and the formula is Cl2. Chlorine 
is a rather more active element than oxygen and combines with 
almost all the other elements to form compounds called chlorides. 
In many cases the combination takes place with the production 
of light and heat, and closely resembles the combustion of sub- 
stances in oxygen. So striking is the resemblance that we call 
the phenomena combustion in chlorine. In fact chlorine reacts 
so vigorously with a number of other substances that they will 
take fire spontaneously when introduced in the gas at ordinary 
temperatures. Among these may be mentioned thin copper foil, 
phosphorus, and powdered antimony. In each case a chloride is 

When completely dry, chlorine is much less active than when 
moist, and carefully dried liquid chlorine is now an article of 
commerce, being sold in strong steel cylinders lined with lead 
or bronze. One of the most important chemical properties of 
chlorine is that it is able to destroy many coloring substances, 
and it is therefore extensively used as a bleaching agent in 
the cotton and paper industries. It bleaches only in the pres- 
ence of water, and is not able to destroy many of the mineral 
dyes nor black tints due to carbon in the form of lamp black. 
Chlorine makes all animal fibers very weak and therefore cannot 
be used to bleach wool or silk. 

It is a very powerful disinfectant and is much used to destroy 

Chlorine Water. As has been mentioned, chlorine is fairly 
soluble in water. The solution, which has the greenish-yellow 
color, the odor, and taste, and bleaching qualities of the gas is 
called chlorine water. A portion of the gas undergoes a chemical 
change with the water, forming compounds which can best be 
discussed later. When exposed to sunlight, the chlorine dis- 
appears, oxygen is given off, and hydrochloric acid is left. The 
equation is 

2H 2 O + 2C1 2 = 4HC1 + 2 

This reaction appears to be just the reverse of that taking 
place in the Deacon Process, but there is this difference that the 


Deacon Process takes place in the gaseous state, while this 
goes on in dilute solution; and under these circumstances the 
hydrochloric acid and oxygen are more stable than chlorine 
and water. 

The action of chlorine upon water takes place slowly in the 
dark but much more rapidly in the light, and the more rapidly 
the stronger the light. The different kinds of light act quite 
differently; the red rays are almost without effect while the 
blue, violet and ultra-violet rays are especially active. This 
action of light is an example of what is known as photo-chemical 

Chlorine Hydrate. When chlorine is passed into ice-cold 
water, there soon separates out a greenish crystalline substance 
which is a compound of water and chlorine having the formula 
Cl2-8H 2 and is called chlorine hydrate. It is a rather unstable 
compound breaking down at room temperature into chlorine and 

Chlorine hydrate is now being prepared on a large scale as one 
of the steps in the process of manufacturing liquid chlorine 
from the impure gas obtained by the electrolysis of sodium 
or potassium chlorides The chlorine is cooled to a low tempera- 
ture and exposed to the action of a spray of cold water. The 
hydrate is formed thus removing the greater part of the chlorine 
from the gaseous mixture. The hydrate decomposes at higher 
temperature into water and chlorine, the latter is dried and is then 
practically pure and ready for liquefaction. 

The Naming of Chlorides. As has been mentioned, the com- 
pounds of other elements with chlorine are usually called chlorides, 
prefixing the name of the other element to form the name of the 
compound, as for example sodium chloride, potassium chloride, 
etc. If, as is not infrequently the case, the same element will 
form more than one compound with chlorine, it becomes necessary 
to distinguish between these substances. Where there are only 
two of these compounds, they are commonly distinguished by 
adding to the name of the element other than the chlorine the 
terminations -ous and -ic. The -ous being used for the compound 
containing the smaller and the -ic for that containing the larger 
amount of chlorine per combining weight of the other element. 
There are two chlorides of iron, the one FeC^ is called ferrous 



chloride, and the other FeCla is called ferric chloride. Similarly 
we have mercurous chloride HgCl and mercuric chloride HgCl 2 . 
In some cases, particularly where there are more than two 
compounds of an element with chlorine, the necessary distinction 
is made by prefixing to the chloride the Greek word signifying 
the number of combining weights of chlorine to a combining 
weight of the other element. For example we have phosphorus 
trichloride, PCls and phosphorus pentachloride, PCU. This same 
system of nomenclature is applied in distinguishing between 
the compounds in other cases in which the same pair of elements 
unite in different proportions to form more than one compound. 
For example, in the case of the oxides we have mercurous oxide, 
Hg 2 O; mercuric oxide, HgO; ferrous oxide, FeO; ferric oxide, 
Fe 2 O3; sulphur dioxide, S0 2 and sulphur trioxide, SOs. 

Hydrogen Chloride. Hydrogen chloride is an exceedingly 
important compound of chlorine. It may be formed by the 
direct combination of the two elements, which can be brought 
about in a number of ways most of which closely resemble those 
which will induce the combination of hydrogen and oxygen. 
For example, a stream of hydrogen will burn in an atmosphere 
of chlorine much as in air except that the color of the flame 
will be a peculiar green instead of the blue of the ordinary flame. 
A colorless gaseous compound is formed which fumes when 
brought in contact with the air. This is hydrogen chloride. 
It is much more readily soluble in water than chlorine, does not 
bleach as chlorine does, but turns blue litmus red. Its solution 
in water does not in any way resemble chlorine water, but is 
highly acid and is identical with hydrochloric acid, of which we 
have had much to say. It has an odor which is disagreeable, but 
entirely different from that of chlorine. It has a very sour taste, 
like all other acids, and in this respect also it differs materially 
from chlorine. In fact hydrogen chloride is unlike chlorine in 
color, odor, taste, solubility, action on colors, and in every other 
property, both physical and chemical. 

Another method for bringing about the combination of hy- 
drogen and chlorine is to mix the gases and pass an electric spark 
through a portion of the mixture. As was the case when this was 
done with a mixture of hydrogen and oxygen, the gases at once 
combine, with a violent explosion. If equal volumes of hydrogen 


and chlorine are used, both of the gases entirely disappear, and 
the volume of the hydrogen chloride produced is equal to the 
sum of the volumes of the separate gases. This shows that the 
combination follows Gay Lussac's law of combining volumes. If 
an excess of either gas is used, that excess is left uncombined. 
The equation for the reaction is 

H 2 + C1 2 = 2HC1 

which shows that one mole of hydrogen, 2.016 parts by weight 
of hydrogen, will combine with one mole of chlorine, 70.92 parts 
by weight of chlorine to form two moles of hydrogen chloride, 
or 2 X 36.468 parts by weight of hydrogen chloride. The compo- 
sition of hydrogen chloride is 1.008 parts by weight of hydrogen 
to 35.46 parts by weight of chlorine, or 2.76 per cent, hydrogen 
and 97.24 per cent, chlorine. 

If the mixed gases be heated to a temperature of between 
240 and 270C. explosion takes place. Hydrogen and oxygen, 
it will be recalled, behave in much the same way except that the 
combination takes place at a higher temperature. 

There is one method for bringing about the combination of 
hydrogen and chlorine which entirely fails in the case of hydrogen 
and oxygen. This is "with light. A mixture of hydrogen and 
chlorine combines very slowly in absolute darkness. If the mix- 
ture be exposed to light, the combination takes place at a rate 
which can be measured and which is proportional to the intensity 
of the light. In ordinary diffused daylight the reaction is com- 
plete in a few days, while if the mixture be exposed to full sunlight 
or to the light of burning magnesium, explosion almost instantly 

Photochemical Action. The action of light in accelerating the 
union of hydrogen and chlorine is another example of photo- 
chemical action. In this case the light certainly does not have 
to impart energy to the mixture, because the combination takes 
place even in the dark and always with a very great decrease 
in the energy so that a large amount of heat is evolved. A 
completely dry mixture of chlorine and hydrogen is insensitive 
to light, so water plays an important part in the process. 

Action of Chlorine on Hydrogen Compounds. Oxygen, it will 
be recalled, was able to act upon many compounds of hydrogen in 



such a way as to produce water. Similarly, chlorine will react 
with hydrogen compounds for the formation of hydrogen chlor- 
ide. We have had an example of this in the case of chlorine and 
water. A still more striking case is that of chlorine and turpen- 
tine. Turpentine is a compound of hydrogen and carbon 
CioHi 6 . If a piece of filter paper be dipped into some slightly 
warmed turpentine and then introduced into a cylinder of 
chlorine (Fig. 20, a), the latter will combine with the hydrogen 
of the turpentine with the production of so much heat that after 
a moment the turpentine bursts into a dark red flame which de- 
posits all the carbon in the form of soot, and forms hydrogen 
chloride which may be identified by its properties. Parafiine 
is composed of compounds of hydrogen and carbon, and when a 
burning paraffine candle (Fig. 20,6) is introduced into chlorine, 
it continues to burn, but with a darkened flame and the formation 

of hydrogen chloride and of soot, because 
under these circumstances chlorine does 
not combine with carbon. 

When natural gas is mixed with chlorine 
and exposed to sunlight, the methane, 
CH 4 , which is the chief constituent of the 
gas, is attacked and hydrogen chloride, 
together with a number of chlorine sub- 
stitution products of methane are formed. 
The first of these products is methyl 
chloride, CH 3 C1, and the final one is carbon tetrachloride, CC1 4 . 
These actions take place slowly and at ordinary temperatures. 

Laboratory and Technical Preparation of Hydrogen Chloride. 
The methods previously given are not suitable for the prepara- 
tion of any considerable quantity of hydrogen chloride. It 
therefore is almost invariably prepared by acting upon common 
salt sodium chloride with sulfuric acid. This action goes on in 
two stages. The first takes place at very moderate temperatures 
and consists in acting upon one mole of sodium chloride with one 
mole of sulfuric acid to form a mole of sodium acid sulfate and 
one of hydrogen chloride as shown in the equation 

NaCl + H 2 SO 4 = NaHS0 4 + HC1 
At a considerably higher temperature the second stage takes 

FIG. 20. 



place and consists in acting upon a mole of sodium chloride with 
a mole of sodium acid sulfate to form a mole of neutral sodium 
sulfate and one of hydrogen chloride, 

Nad + NaHS0 4 = Na 2 SO 4 + HC1 

These two equations may be added to- 
gether to obtain one which will represent 
the complete process 

2NaCl + H 2 S0 4 = Na 2 SO 4 + 2HC1 

The operation is carried out in the 
laboratory in a glass flask, (Fig. 21) and 
the temperature cannot be raised high 
enough to carry it beyond the first stage. 
On a manufacturing scale the process is 
carried through the first stage in a large 
iron pan, A, heated from below (Fig. 22) , 
covered with a brickwork dome; and 
finished in a fire clay muffle, B. Both 
dome and muffle are connected by a flue 
with brick towers filled with coke over 
which a small stream of water is kept 
trickling. The hydrochloric acid vapors pass up through the 
towers against the stream of water and are absorbed. Almost 
saturated hydrochloric acid runs out from the bottom of the first 

FIG. 21. 

Pan Purifiers 

FIG. 22. 


The reactions just given are reversible and would be incomplete 
were it not for the fact that the hydrogen chloride is very volatile 
and hence readily distilled from the mixture. This, of course, de- 
creases the concentration of the hydrogen chloride, and more is 
at once formed to take its place. This will continue until the 


salt and sulf uric acid have been completely transformed, provided 
that the hydrogen chloride is carried off as it is formed. 

The action takes place not because the hydrochloric acid is a 
weaker acid than sulfuric acid it is much stronger but because 
the hydrochloric acid is much more volatile than sulfuric acid. 
In fact sulfuric acid may be replaced by any acid, however weak, 
which is much less volatile than hydrochloric acid and which is 
not a strong enough oxidizing agent to attack the latter. Phos- 
phoric acid will do very well. 

First General Method for the Preparation of Acids. In the 
cases given above we are dealing with examples of a general 
method for the preparation of acids from their salts. This 
may be called the first general method for the preparation of 
acids and may be stated as follows. Treat a salt of the acid 
desired with a less volatile acid and distil. It is applicable to such 
acids as will stand distillation without decomposition, and of 
course the added acid must not only be less volatile than the acid 
desired but must also be without action upon this acid. 

Properties of Hydrogen Chloride. A number of properties of 
hydrogen chloride have already been given; in addition we may 
state that its density is 0.00164 and molecular weight, 36.47. Its 
critical temperature is 51.4C. and its critical pressure 81.6 at- 
mospheres. It can then be converted into a liquid by pressure 
alone at ordinary temperatures. At 22C. the pressure neces- 
sary is 46 atmospheres. The boiling-point of the liquid hydrogen 
chloride is -83C. and the freezing-point is -110C. The 
liquid hydrogen chloride is a colorless almost inactive sub- 
stance which is without action upon metals or blue litmus, and 
in general exhibits very little of the great chemical activity which 
is shown by its solution in water. This difference is of great 
import as we shall soon see. 

As has been mentioned, hydrogen chloride is very soluble in 
'water. At its freezing-point one volume of water will absorb 
about 525 volumes of hydrogen chloride. At the same time a 
great amount of heat is developed which points to the formation 
of a compound with the water. This is indicated too by the 
fact that hydrogen chloride does not follow Henry's law of solu- 
bility of gases which states that provided no chemical reaction 
takes place between the gas and the solvent, the mass of the gas dis- 



solved by a given mass of the solvent is directly proportional to the 
pressure. In the case of hydrogen chloride and water, Henry's 
law is far from describing the behavior. In fact the mass of the 
hydrogen chloride dissolved is but slightly altered by large 
changes in pressure. This would indicate that the hydrogen 
chloride undergoes some sort of a change upon passing into 
solution in water. As a matter of fact by cooling very concen- 
trated solutions of hydrogen chloride to low temperatures, three 
crystalline compounds of hydrogen chloride and water can be 
obtained. These are the mono-, di- and tri-hydrates, and have 
the formulas HC1-H 2 O, HC1-2H 2 0, and HC1-3H 2 0. 

Constant Boiling Hydrochloric Acid. Hydrogen chloride 
boils at 83C. and water at 100C. so one would naturally 


expect that a solution of hydrogen chloride would boil at tempera- 
tures between these two points. This is true, however, only 
of solutions containing more than 40 per cent, of hydrogen 
chloride. Weaker solutions boil at temperatures above 100C. 
and a 20 per cent, solution boils at 110C. and is the highest 
boiling and consequently the least volatile mixture of hydrogen 
chloride and water. If one distils a hydrochloric acid solution of 
less than 20 per cent, both water and hydrogen chloride come off, 
but relatively more of the former, and the residue becomes richer 
and richer in hydrogen chloride. After this process has gone on 
for some time, the residue finally reaches 20 per cent, hydrogen 
chloride which is the least volatile of the hydrogen chloride- 


water mixtures, and boils at 110C. After this has been reached, 
the mixture distils unchanged in composition and boiling-point. 

If on the other hand we start to distil a hydrochloric acid solu- 
tion containing more than 20 per cent, hydrogen chloride, 
this substance will come off with the water but in such quantities 
that the distillate contains a higher percentage of hydrogen 
chloride than the original mixture, leaving a residue which con- 
tains a smaller .percentage of hydrogen chloride and boils at a 
higher and higher temperature until its boiling-point becomes 
110C. and its composition 20 per cent, hydrogen chloride. 
After this it boils constantly with unchanged composition. From 
the above it follows that no matter what the original composition 
of the hydrochloric acid, after it has boiled for sufficient time the 
residue will be a 20 per cent, acid boiling at 110C. 

The relationship between the composition and boiling-point for 
hydrochloric acid is shown in Fig. 23. 

This constant boiling hydrochloric acid has often been mis- 
taken for a chemical compound, but that this is not the case is 
shown by the fact that its composition changes with the pressure. 
When the boiling is carried out under a pressure of 2.5 atmos- 
pheres, the composition is 18 per cent, hydrogen chloride, while 
at 0.066 atmospheres it is 23 per cent, hydrogen chloride. 


The Ionic Theory, or the Theory of Electrolytic Dissociation 
as it is often called, is one of the most important tools of modern 
chemistry, and this chapter will be devoted to its development. 
As should always be the case, the facts which it was devised to 
explain will be given first and should be fully appreciated by the 
student, and then the theory learned not as a fact, but as a pos- 
sible explanation for the facts. 

Acids. Mention has been frequently made in what has gone 
before of acids; hydrochloric, sulfuric and phosphoric acids 
are typical examples of this very important class of compounds. 
The acids comprise a large group of substances of very different 
compositions which have certain well marked properties in com- 
mon. For example, they all contain hydrogen as one of their 
essential constituents, and when in solution they all taste sour, 
redden litmus, and evolve hydrogen gas when brought in contact 
with magnesium or zinc. All acids are hydrogen compounds, 
but many hydrogen compounds are not acids. Water, turpen- 
tine, sugar, olive oil, and many other substances contain hydro- 
gen but will not evolve hydrogen when brought in contact with 
magnesium or zinc at ordinary temperatures, neither do they 
taste sour, nor redden litmus, and so are not acids. The hydro- 
gen, then, of acids must be in some kind of a special condition, 
different from that of the hydrogen of other compounds, and since 
this is the only constituent which acids have in common, as is 
shown by the appended list of acids and their formulas, the com- 
mon properties of acids must be ascribed to the hydrogen in this 
special condition. 


Hydrochloric acid, HC1 Nitric acid, HNOs 

Hydrobromic acid, HBr Phosphoric acid, H 3 PO 4 

Hydriodic acid, HI Oxalic acid, H 2 C 2 O4 

Sulfuric acid, H 2 SO 4 Acetic acid, HC a HjO a 



The acids show their peculiar properties only when dissolved 
in water and a few other solvents. When in a perfectly pure 
state or dissolved in most solvents, they are as indifferent as pure 
hydrogen chloride and neither redden litmus nor evolve hydrogen 
with magnesium. Of course, their taste in the entire absence 
of water cannot be determined. The water has then some- 
thing to do with the transformation of the hydrogen of the acids 
into its special state. In this connection it will be recalled that 
there is evidence that hydrogen chloride undergoes a chemical 
change when it passes into solution in water. 

Bases. Bases are another rather large class of substances of 
varied composition which like acids, have, when dissolved in 
water, certain properties in common. Their solutions all have 
a peculiar alkaline taste, a soapy feeling, and turn red litmus 
blue. As will be seen from the partial list of bases given below, 
they all contain hydrogen and oxygen in the proportion of one 
atomic weight of the one to an atomic weight of the other. 
We express this by saying that they contain the hydroxyl group, 
OH, and ascribe their common properties to this group. 


Sodium hydroxide, NaOH Calcium hydroxide, Ca(OH)2 

Potassium hydroxide, KOH Barium hydroxide, Ba(OH) 2 

Ammonium hydroxide, NH 4 OH Strontium hydroxide, Sr(OH) 2 

Like the acids, these bases show their common properties only 
when dissolved in water and a few other solvents, and since there 
are many compounds known which contain hydroxyl and yet 
are not bases, we ascribe the characteristic properties of bases 
in water solution to the hydroxyl in a peculiar condition, similar 
to that of the hydrogen from acids in such solutions. 

Neutralization. When a solution of hydrochloric acid is 
added to a solution of sodium hydroxide in just the proper pro- 
portions, both the characteristic properties of the acid and those 
of the base completely disappear, and the solution which is left 
behind neither turns blue litmus red nor red litmus blue. It has 
neither a sour nor an alkaline taste, has no soapy feeling and does 
not evolve hydrogen at ordinary temperatures with zinc or mag- 
nesium. Its taste is a pure salty one and in fact it is in every 
way identical with a solution of common salt, sodium chloride, 


in water. If the solution be evaporated, nothing but water 
passes off, and nothing but sodium chloride is left behind. What 
then has become of the hydroxyl of the base and the hydrogen 
of the acid? Obviously they must have combined to form water, 
leaving the sodium and chlorine to form sodium chloride. 
The reaction may be represented by the equation 

NaOH + HC1 = H 2 O + NaCl 

We say that the hydrochloric acid and sodium hydroxide have 
neutralized each other and that water and a solution of sodium 
chloride have been formed. Other acids and bases behave in 
much the same way. They neutralize each other and leave 
solutions of the corresponding salts. For example, potassium 
hydroxide and nitric acid give water and a neutral solution of 
potassium nitrate, 

KOH+HNO 3 = H 2 O+KNO 3 
Sulfuric acid gives sulfates and water 

2NaOH+H 2 SO 4 = 2H 2 O+Na 2 SO 4 

Reacting Ratio in Neutralization. By inspection of the 
above equations and writing others, it is very easy to satisfy 
one's self that a quantity of a base which contains 17.008 grm. 
of hydroxyl will just neutralize a quantity of an acid which 
contains 1.008 grm. of hydrogen in the peculiar acid condition. 
Since it is a very easy matter to tell with the aid of litmus just 
when a solution is neutral, the fact mentioned above is made the 
basis of a simple and much used method for the quantitative 
determination of acids and bases. The essentials of the method 
are as follows : Such a quantity of a base, say sodium hydroxide, 
as will contain 17.008 grm. of hydroxyl is dissolved in water 
and the solution made up to just a liter. Each cubic centimeter 
will then contain 0.017008 grm. of hydroxyl and will neutralize 
that quantity of an acid which will contain 0.001008 grm. of 
hydrogen, that is- to say, 0.03647 grm. of hydrogen chloride for 
example. Such a solution is called a normal solution of a base. 
To use this in determining the amount of an acid, say hydrogen 
chloride, in a solution of unknown strength one would weigh 



out a few grams of the acid, add a drop of litmus solution and 
then run in from a graduated vessel called a -burette (Fig. 24), 
enough of the normal hydroxide solution so that the litmus is 
just on the point of turning blue, and in fact is neither blue 
nor red. The number of cubic centimeters used multiplied by 
0.03647 gives the weight of hydrogen chloride in the acid solution 

taken. Suppose that we weighed 
out 10 grm. of the acid and it re- 
quired 20 c.c. of the normal base to 
neutralize it. The acid contained 
in this 10 grm. is 

20X0.03647 = 0.7294 or 7.294 
per cent. 

Similarly with the aid of a normal 
acid, i.e., one containing 1.008 grm. 
of hydrogen in the acid condition 
per liter, one may determine very 
easily the quantity of base in a 
solution of unknown strength. 

Methods like the above are in 
very frequent use in quantitative 
chemical analysis, and are called 
volumetric methods. 

Acids, Bases, and Salts have 
two Sets of Properties. Every 
acid has when in solution in water, 
w/w two independent sets of properties, 
one common to all acids and evi- 
dently belonging to the hydrogen 

of the acid, and the other peculiar to that acid, being shown by 
no other acid, and evidently belonging to the rest of the acid 
other than the hydrogen. For example, hydrochloric acid is 
like all other acids in that it is sour, reddens litmus and gives 
hydrogen with zinc, but differs from all other acids in that 
it will give sodium chloride with sodium hydroxide, and also 
will yield a white curdy precipitate with silver nitrate ^hich 
precipitate is silver chloride. A few other acids give somewhat 
similar precipitates with silver salts, but these all differ sufficiently 


FIG. 24. 


in their properties from the silver chloride so that they may be 
positively distinguished from the latter. 

Bases too have two independent sets of properties in water 
solutions, the one common to all bases and evidently due to the 
hydroxyl, and the other peculiar to the particular base that is 
being worked with and belonging to the rest of the base other 
than the hydroxyl. For example, barium hydroxide shows all 
the characteristic properties of bases, turning litmus blue, etc., 
and also gives with sulfuric acid and all sulfates a white precipi- 
tate of barium sulfate, BaSO 4 , which is different in its properties 
from all other substances. No other base has just this property. 
When one of these acids with its two sets of properties is neutral- 
ized by a base with its two sets, the properties common to all 
acids together with those common to all bases completely dis- 
appear leaving behind that set of properties of the acid which 
was peculiar to that particular acid, and the set of properties of 
the base which was peculiar to that base. So the salt solution 
remaining after the neutralization of an acid by a base has two 
sets of properties. 

It follows from what has just been said that solutions of hydro- 
chloric acid and all soluble salts of hydrochloric acid should give 
the white curdy precipitate of silver chloride upon the addition of 
silver nitrate. This is exactly what has been found to be the 
case. Similarly barium salts formed by the neutralization of 
this base by acids should and do give precipitates of barium 
sulfate upon the addition of sulfuric acid or sulfates. 

When these same bases, acids and salts are in the pure state 
or even in solution in most solvents other than water, they do not 
exhibit these two sets of independent properties, but each has 
only one. An explanation for these and for the succeeding sets 
of facts will soon be given as the theory of electrolytic dissociation. 

Abnormal Lowering of the Freezing-point. In connection 
with the discussion of hydrogen peroxide it was brought out that 
a gram mole of a normal substance lowers the freezing-point of 
1,000 grm. of water 1.85C. and that this may be made the basis 
of a method for the determination of molecular weight. In fact 
we got the molecular weight of hydrogen peroxide in this way. 
The law of the molecular lowering of the freezing point is more 
exact for dilute solutions than for concentrated ones and we will 


be more accurate if we state it as follows : A gram mole of a nor- 
mal substance lowers the freezing point of 1000 kgrm. of water 

The acids, bases and salts are the exceptional substances 
which do not lower the freezing-point of 1,000 kgrm. of water 
0.00185C. per gram mole, but always produce a greater lowering. 
In many cases the lowering is nearly twice what we would expect, 
in others something like three times, and sometimes even more 
than three times the 0.00185C. per gram mole dissolved in 1,000 
kgrm. of water. 

For example, hydrogen chloride is a gaseous substance and 
its molecular weight can be found from its gaseous density to be 
36.47. If one dissolves 36.47 grm. of hydrogen chloride in 1,000 
kgrm. of water, the freezing-point of the water is lowered almost 
twice 0.00185C. A gram mole of sodium or potassium chloride 
or of any other salt formed by the union of one atomic weight 
of a metal with one atomic weight of chlorine produces about the 
same lowering as a gram mole of hydrogen chloride when dis- 
solved in 1,000 kgrm. of water or about twice 0.00185. Such 
salts as calcium or barium chloride, CaCU or BaCl 2 , which con- 
tain two atomic weights of chlorine to one atomic weight of the 
metal, lower the freezing-point nearly three times 0.00185 per 
mole of salt in 1,000 kgrm. of water. From these facts it looks 
as though the lowering for salts were 0.00185C. per gram atomic 
weight of each element instead of per gram mole. That this is 
not the case is shown by the fact that nitric acid, HNOs, potas- 
sium nitrate, KNO 3 , and sodium nitrate, NaN0 3 , only lower the 
freezing-point of 1,000 kgrm. of water approximately twice 
0.00185 per gram mole of acid or salt. In more concentrated 
solutions these acids and salts give lowerings which are abnormal 
but not as much so as in the dilute solutions; however, even in 
the most dilute solutions the abnormality is never greater than 
the figures just given. These facts are very significant and 
important in connection with the theory soon to be developed. 

Electrolytes. Metals are called conductors of the first class. 
As is well known, they are not altered in any way by the current 
except that their temperature is more or less raised. Some 
solutions are also conductors of electricity and are called con- 
ductors of the second class because they are not only become 



heated, but also are invariably decomposed during the passage of 
the current provided the circuit consists in part of conductors of 
the first class as is usually the case. Such solutions are called 
electrolytes and are formed only when acids, bases, and salts 
are dissolved in water or in the few other solvents in which they 
each show their two independent sets of properties. The de- 
composition takes place at points where the electric current 
enters and leaves the salt solution. The terminals of the metallic 
portion of the circuit are called electrodes and are distinguished 
from each other by calling that electrode where the electricity 

Chlorine- bi 

FIG. 25. 

enters the solution the anode and that where it leaves it for 
the metal, the cathode. During the passage of the current, 
one kind of substance appears at the cathode and an entirely 
different substance at the anode. When hydrogen chloride solu- 
tion is electrolyzed Fig. 25, hydrogen appears at the cathode and 
chlorine at the anode. The two substances appearing at the 
electrodes are in general the two substances which seem to give to 
the salt solutions their two sets of independent properties. 

A fact which is closely connected with those given above, and 
which is still more significant is that, when a current of electricity 
is passed through a salt solution, the two components of the salt 
which act chemically independently of one another may actually 
be seen to move slowly in opposite directions, the one toward the 
cathode and the other toward the anode, and at quite different 
rates. Especial precautions must be taken to bring this out, 
but the experiments are rather easily performed and are quite 
definite in their results. 


In the case of hydrochloric acid solutions, the chlorine moves 
toward the anode while the hydrogen travels toward the cathode, 
and moves nearly five times as rapidly as the chlorine. 

The Law of Faraday. The English scientist Faraday in the 
year 1833 carefully investigated the phenomena occurring at the 
electrodes and devised the names which we have been using 
such as electrolysis, electrolyte, electrode, anode, cathode, and in 
addition he called the substances which travel toward the elec- 
trodes ions. Those which move toward the cathode being the 
cations and those which move toward the anode being the anions. 
He also discovered facts which are described by the folio-wing 
law known as the Law of Faraday. 

Electricity moves in electrolytes only with the simultaneous move- 
ment of their chemically independent components, and is accom- 
panied by the deposition of two different substances upon the elec- 
trodes, the quantity of electricity required to liberate a gram atomic 
weight of any substance being 96,500 coulombs or some integral 
multiple of this. 

Summary. We have now discussed the principal groups of 
facts which the theory of electrolytic dissociation was devised to 
explain, but before proceeding to the theory itself a brief summary 
of these facts may be of use. 

First. Salts in solution in water showitwo independent sets 
of chemical properties; all salts having a common component 
have one set of properties in common. The term salt as used 
here includes acids and bases as well as salts in the ordinary sense. 

Second. Salts in solution in water lower the freezing-point of 
the water more than 0.00185C. per gram mole per 1,000 kgrm. 
of water; some nearly twice, some three times, and some even 
more than three times the normal lowering. 

Third. Salts in solution are electrolytes and the electricity 
passes through the solution only with the simultaneous move- 
ment of the components of the salts in opposite directions and at 
quite different rates. In addition two different substances are 
liberated at the electrodes. One gram atomic weight of each 
substance is liberated by the passage of 96,500 coulombs or some 
rational multiple of this. 

All these are well established facts. The theory which is about 
to be given is simply an imagined cause or explanation for these 


facts and is in many respects on a par with the atomic theory. 
It is very useful in correlating a great many facts of chemistry 
and has been very successful in predicting the discovery of new 
facts, and hence is worthy of considerable attention. 

The Theory of Electrolytic Dissociation. To explain the facts 
outlined above, certain assumptions are made concerning the 
nature of salts in solution. The first one is that when salts are 
dissolved in water, they are immediately decomposed or dissoci- 
ated into at least two different substances which are chemically 
separate and distinct from each other and from the original salt. 
These new substances are assumed to be the chemically inde- 
pendently acting components of the salts and are called ions, from 
the Greek word ov meaning wander, because they move in oppo- 
site directions toward the electrodes. That which goes toward 
the anode is known as the anion and that toward the cathode as 
the cation. 

Each gram mole of salt which dissociates is assumed to give 
rise to at least 2 gram moles of these new substances, the ions, 
and each gram mole of an ion, or for brevity "gram ion," is 
assumed to act in lowering the freezing-point of water like a gram 
mole of any other substance. 

To account for the electrical properties of salt solutions, it is 
assumed that these new substances, the ions, are charged with 
electricity, the cations with positive and the anions with negative 
electricity; the charges amounting to 96,500 coulombs or a 
integral multiple per gram ion. Since salt solutions are electric- 
ally neutral, it must be assumed that the + charges upon the 
cations just equal the charges upon the anions. 

The explanation of the first group of facts, p. 108, by the theory 
is very simple. We have only to say that each of the two inde- 
pendent sets of properties belongs to one of the new substances 
or ions formed by the dissociation of the salt, and since these ions 
are assumed to be uncombined, each set of properties is of course 
independent of the other component of the salt. The chlorine ion 
is assumed to be identical whether it comes from hydrochloric 
acid, HC1, sodium chloride, NaCl, potassium chloride, KC1, or 
any other chloride, and hence all chlorides in solution show the 
properties of chlorine as ion no matter what the other ion may be. 
The acid properties are ascribed to hydrogen as ion and since this 


is assumed to be present in all acids in solution they all show the 
common properties of acids irrespective of the nature of the 
other ion. 

The abnormal lowering of the freezing-point is explained by 
the assumption that each gram mole of salt when it dissociates 
produces two or more gram moles of ions, or gram ions as they 
are called. A gram mole of hydrogen chloride, it will be recalled, 
lowers the freezing-point of 1,000 kgrm. of water practically twice 
0.00185C. This gram mole of hydrogen chloride is supposed, 
by the theory, to decompose into a gram mole of hydrogen as ion 
which lowers the freezing-point of the 1,000 kgrm. of water 
0.00185 and a gram mole of chlorine as ion which also lowers the 
freezing-point another 0.00185 or 0.00370 for the two together. 

It must be noted that a gram mole of hydrogen as ion is 
assumed to differ from a gram mole of hydrogen in that it has 
only one atomic weight of the element instead of two per mole 
and carries 96,500 coulombs of electricity. The same kind of 
a relation exists between the gram ion and the gram mole of 

The fact that a gram mole of calcium chloride, CaCl2, lowers 
the freezing-point of 1,000 kgrm. of water nearly three times 
0.00185 is easily explained because the theory assumes that it 
breaks up into a gram mole of calcium as ion arid two gram moles 
of chlorine as ion. Each lowers the freezing-point 0.00185 or 
three times this for the three of them. Sulfuric acid acts as 
though it breaks up into two gram moles of hydrogen as ion and 
a gram mole of sulfate, SO 4, and lowers the freezing-point of the 
standard 1,000 kgrm. of water nearly 3 X 0.00185 as would be 

The movement of the components of the salts, the one toward 
the anode and the other toward the cathode and at quite different 
rates can be easily explained with the aid of the assumptions 
made above. The ions are supposed to be charged with elec- 
tricity, the cations positively and the anions negatively. The 
anode is charged positively and the cathode negatively, just the 
opposite from the ions. Now from the well-known fact that 
electricities of opposite signs attract each other, it is easy to see 
why the cations should go toward the cathode and the anions 
toward the anode. Since the water is supposed to have broken 


up the salt before it was electrolyzed, no further explanation is 
necessary, for the components traveling in opposite directions 
and at different rates, since there is nothing to hinder them from 
doing so because they are no longer chemically combined with 
each other. The law of Faraday is easily explained by assum- 
ing that each gram ion carries 96,500 coulombs or some integral 
multiple of this which must be neutralized by electricity from 
the electrode before the ion is deposited at the electrode and 
hence each gram ion will require 96,500 coulombs or some 
integral multiple of this before it can be liberated. 

The fact that only solutions of salts are conductors of elec- 
tricity is accounted for by saying that it is only in such solution 
that ions exist to carry the current. This is checked up by the 
fact that non-conducting solutions give normal lowerings of the 
freezing-point and show only one set of properties. 

Even these salts, of which we have had so much to say, when 
dissolved in some solvents other than water are non-conductors 
and then they do not show the two sets of properties and act 
normally with respect to the freezing-point lowering. So there 
must be something of a connection between these three sets of 
facts, and the theory of electrolytic dissociation is the best guess 
concerning this which has been made. 

In terms of the atomic theory the ions become charged atoms 
or groups of atoms, but a very good way to look at them is to 
regard them as substances which are charged with electricity and 
in a peculiarly reactive condition. 

The ions are represented by the symbols of the element or 
elements going to make them up, together with as many + or 
signs as they seem to carry multiples of 96,500 coulombs of + or 
electricity. The following list gives the names and symbols of 
a few of the more common ions. 


Hydrogen H+ Chlorine Cl~ 

Sodium Na+ Bromine Br~ 

Cotassium K + Hydroxyl OH~ 

Silver . . Ag+ Nitrate NO 3 ~ 

Palcium Ca + + Sulf ate SO 4 ~ ~ 

Barium Ba ++ Phosphate PO 4 

Ammonium . . . . NH4 + 


In terms of this theory a salt is any substance which when 
dissolved in water yields ions. 

An acid is a salt which when dissolved in water gives hydrogen 
as one of its ions. 

A base is a salt which when dissolved in water gives hydroxyl 
as one of its ions. 

Neutralization of Acids and Bases. When we come to consider 
the neutralization of acids and bases in the light of the theory, we 
reach the rather surprising conclusion that in the process only 
water is formed. The salt in the narrow sense, is not formed 
since its cation and anion existed as such in the solution of the 
base and acid before they were brought together and continue 
to do so after the process is finished. 

To illustrate, let us give the equations for a few reactions of 
this kind. 

Na+ + OH~ + H+ + Or = H 2 O + Na+ + Cl" 

K+ + OH- + H+ + cr = H 2 o + K+ + cr 

Na+ + OH- + H+ + NO 3 - = H 2 O -f Na+ + NOr 
K+ + OH- + H+ + N0 3 - = H 2 + K+ + NOr 
2Na+ + 20H~ + 2H+ + S0 4 " "= 2H 2 + 2Na+ + SO 4 ~ ' 
Ca++ + 20H- + 2H+ -f 2C1~ = 2H 2 + Ca++ + 2C1~ 

In every case water is formed and water only. The cation of 
the base and the anion of the acid are left unaltered. This 
explains why these salt solutions show the properties which were 
peculiar to the particular base and acid from which they were 

Heat of Neutralization. Whenever any of the acids and 
bases given in the above list neutralize each other there is always 
57 Kj. of heat evolved per mole of water produced, and this is 
true of all other strong acids and bases as well. This fact is very 
easy to explain on the basis of the electrolytic dissociation theory. 
In the light of this theory, water is the sole product of these re- 
actions and is always formed from the same substances, hydrogen 
as ion and hydroxyl as ion, and therefore each reaction should 
show the same heat effect; and thus 57 Kj. or 13,600 cal.is the heat 
of formation of water from hydrogen and hydroxyl as ions. If 


the salts had really been formed, the heat effect should have been 
different in each case. 

Action of Acids. In our previous discussion of acids, certain 
points of similarity, their taste, action on litmus, metals, etc., 
were made much of and ascribed to their common constituent 
hydrogen, in what we termed the ionic state. In other respects 
acids show great similarity. Quantities of different acids which 
contain equal amounts of ionizable hydrogen are equivalent to 
each other in several ways. They will, for example, neutralize 
the same quantity of any given base, say sodium hydroxide, and 
also when brought in contact with an excess of metal, like zinc, 
evolve equal quantities of hydrogen. Such quantities of acid 
as contain equal weights of ionizable hydrogen are appropriately 
called equivalent quantities, and that weight of any given acid 
which contains 1.008 grm. of ionizable hydrogen is called a gram 
equivalent of that acid. It will be recalled that a gram equiva- 
lent of an acid in 1,000 c.c. of solution is called a normal solution 
of the acid. 

In some respects acids show striking points of dissimilarity. 
If pieces of zinc of equal area are introduced into equivalent 
solutions of hydrochloric, sulfuric, and acetic acids, hydrogen 
will be evolved in each case, but at very different rates in the 
different acids. If normal acids are used, less than 1 c.c. of 
hydrogen will be liberated from the acetic acid, and 65 c.c. 
from the sulfuric acid during the time required for the liberation 
of 100 c.c. from the hydrochloric acid. If equivalent quantities 
of the acids and an excess of zinc are used, ultimately, equal 
quantities of hydrogen will be liberated from each. In order to 
insure the success of the experiment indicated above, the pieces 
of zinc, before use should be treated with a dilute solution of 
copper sulfate. Copper will be deposited on the zinc and will 
tend to overcome the disturbing influence of local impurities in 
the zinc and secure a uniform evolution of hydrogen. 

If we examine other cases in which these acids take part in 
reactions at measurable rates, we discover that their rates stand 
in the same order as in their action on zinc. 

Much the same relationship exists between the lowering of the 
freezing-points of these acids. One gram mole of hydrochloric 
acid, it will be recalled, lowers the freezing-point of 1,000 grm. 


of water nearly twice the normal lowering of 1.85 per gram 
mole; while a gram mole of acetic acid per 1,000 grm. of water 
lowers it only a trifle more than 1.85. 

Normal solutions of these acids show very much these same 
differences in their power to conduct electricity. The solution 
of hydrochloric acid is a good conductor; sulfuric acid something 
like two-thirds as good; while acetic acid, although still a con- 
ductor, is a poor one. 

Because of these peculiarities in their actions, everyone calls 
hydrochloric and sulfuric acids strong acids, and acetic acid a 
weak acid. The hydrochloric acid is considered to be somewhat 
stronger than sulfuric. 

Explanation. It is a very simple matter to explain the facts 
outlined above with the aid of the ionic theory. We have only 
to add to our previous assumptions, that all acids break up or 
ionize so as to give hydrogen as ion, and that this hydrogen as ion 
has the characteristic acid properties, the further assumption 
that the various acids differ among themselves in the extent to 
which they break up, or in their degree of ionization. Hydro- 
chloric acid is assumed to break up or ionize so far that in normal 
solution the greater part of the substance has been changed into 
hydrogen and chlorine as ions, while under the same conditions 
only a very small part of the acetic acid is supposed to be ionized; 
while sulfuric acid stands between the other two. Naturally, if 
hydrogen as ion is the hydrogen which acts acid, that solution 
which contains the greatest concentration of hydrogen present 
actually as ion will be the most intensely acid. This, from what 
has been said, would be the hydrochloric acid, then would come 
sulfuric and last the acetic far behind the other two. 

We have learned that the rate at which any given substance 
will react, other things being equal, depends upon the concen- 
tration, and further that in terms of the ionic theory the reaction 
between an acid and zinc is between hydrogen as ion and the zinc. 
From all that has been said above, then, it follows that the rate 
of evolution of hydrogen gas from these three acids should stand 
in the order, hydrochloric, sulfuric, and acetic acids. This was 
the order found by experiment. 

But if the acids differ so much in their degree of dissociation or 
ionization, how does it happen that equivalent quantities will 


neutralize the same amount of sodium hydroxide or any other 
base, and also ultimately evolve the same quantity of hydrogen 
when acted upon by an excess of a metal? To explain this we 
have only to assume that the ionization of an acid is a reversible 
process, and that as the hydrogen as ion is used up by combin- 
ing with hydroxyl or reacting with the metal, more of the 
undissociated acid will break up into ions, and that this 
process will continue until the undissociated acid is completely 
used up. 

The total acid hydrogen of an acid is called the ionizable hydro- 
gen and that fraction of it which is present at any one instant in 
the state of the free ion, is called the actual ion, while that which 
is undissociated, but capable of becoming ionized, is called the 
potential ion. The ionizable hydrogen is the sum of the actual 
and the potential ions. Equivalent quantities of these three 
acids then contain the same amount of ionizable hydrogen, but 
differ in the relative proportions of the actual and potential 
ions; hydrogen chloride having the most actual and acetic the 

The explanation of the difference in the lowering of the freezing- 
point is so obvious that it need not be gone into here. 

The ability of a gram mole of an electrolyte to conduct elec- 
tricity is called its molecular conductivity. It may be de- 
termined by ascertaining the number of amperes of electricity 
which will flow through a conductivity cell when the whole of 
the solution containing a gram mole of the solute is placed 
between electrodes 1 cm. apart and a potential of 1 volt is 

As a rule the molecular conductivity of a salt increases as the 
volume of the solution in which one gram mole is dissolved is 
increased, i.e., as the solution is diluted. With the great majority 
of electrolytes, the molecular conductivity increases with the 
dilution up to a certain point, and then runs along without any 
material increase as far as the measurements can be carried 
in other words, it reaches a maximum. 

A natural explanation for this is that the molecular con- 
ductivity of a solution is proportional to the fraction of the gram 
mole which is present as ions or the degree of ionization or disso- 
ciation as it is called and to the speed with which the ions move. 



Assuming that the latter is constant as the solution is diluted, it 
follows that the degree of ionization of the electrolyte increases 
with the dilution, and that when the conductivity has reached 
its maximum, the salt is entirely ionized or completely disso- 
ciated. The degree of dissociation at any given dilution, can be 
found by dividing the molecular conductivity at that dilution by 
the maximum molecular conductivity at very great dilution. 
The results given in the following table have been obtained in 
this way. Dilution as used here means the number of liters of 
the solution which contains one gram mole of the acid. It will 
be seen from this table that even at moderate dilutions, the first 
three acids are almost completely dissociated and are about as 
strong as possible. The sulfuric acid is distinctly weaker, but 
quickly gains in strength. The hydrofluoric and acetic acids are 
much weaker, but grow rapidly stronger. So the more dilute 
the acids the stronger they become in the sense that they are 
more largely dissociated. That acid is the strongest which is most 
easily broken up into its ions. No adequate reason has been 
offered as to why hydrochloric acid is so highly and acetic acid so 
slightly dissociated. 






H 2 SO4i 




























Bases, like acids, differ widely in the degree of their dissocia- 
tion at moderate dilutions. Sodium and potassium hydroxides 
being highly dissociated. Calcium and barium hydroxides have 
about the same degree of dissociation as sulfuric acid at any given 
dilution, while ammonium hydroxide is weak like acetic acid. 

Salts, in the common sense of the term, are with few exceptions 
highly dissociated, and that too whether they are salts of weak 
or of strong acids. Sodium acetate, for example, at any given 
dilution is only a little less dissociated than sodium chloride. 

The sulfuric acid solutions are equivalent and not molecular solutions. 


Hydrogen Ion as An Oxidizing Agent. When an acid acts upon a 
metal and hydrogen is given off, the reaction seems to be between 
the hydrogen ion and the metal. For example, the equations 

Zn + 2H++ 2CT = Zn+++ 2CT-f H 2 
Zn + 2H++ S0r~= Zn++-f SO 4 ~+ H 2 

indicate that neither the chlorine nor the sulfate ion is changed, 
while the zinc is transformed from the metal to the ion and the 
hydrogen from the ion to the free element. This change of the 
zinc from the metal to the ion is considered oxidation and the 
hydrogen ion is the oxidizing agent. This type of oxidation must 
be kept in mind until a full definition of oxidation and reduction 
can be offered. 


General. Neither oxygen and chlorine nor oxygen, hydrogen 
and chlorine combine directly, but by more or less roundabout 
methods three oxides and four acids containing oxygen, oxy acids 
as they are called, can be prepared. The following table gives 
the names and formulas of the acids, the sodium salts of these 
acids, and the oxides. 


HC1 Hydrochloric acid NaCl Sodium chloride 

HC1O Hypochlorous acid NaCIO Sodium hypochlorite 

HC1O 2 Chlorous acid NaClO 2 Sodium chlorite 

HC1O 3 Chloric acid' NaClOs Sodium chlorate 

HC1O 4 Perchloric acid NaClO4 Sodium perchlorate 


C^O, Chlorine monoxide or hypochlorous anhydride. 
C1O 2 , Chlorine dioxide or chlorine peroxide. 
C1 2 O 7 , Chlorine heptoxide or perchloric anhydride. 

These oxyacids and their salts furnish excellent illustrations of 
the law of multiple proportions. 

Nomenclature of Acids and Salts. The names of acids 
formed by the union of hydrogen with one other element and 
which consequently do not contain oxygen are formed by taking 
the characteristic part of the name of the element other than 
hydrogen and adding the prefix hydro- and the suffix -ic. For 
example hydro-chlor-ic acid. Salts of such acids are named by 
combining the name of the metal which replaces the hydrogen 
of the acid with another word formed by dropping the prefix 
hydro- from the name of the acid and replacing the -ic by -ide. 
For example, sodium chloride. 

In naming the oxyacids the best known is usually designated 
by adding to the characteristic part of the name of the element 



other than hydrogen and oxygen the suffix -ic. For example, 
HC10 3 is called chloric acid. The names of the salts of such acids 
end in -ate, as sodium chlorate, NaClOs. 

If there is an acid containing more oxygen than the -ic acid, it 
is distinguished by adding to the name of the -ic acid the prefix 
per- and similarly for the salt as perchloric acid, HC1O 4 , and 
sodium perchlorate. That acid containing the next smaller 
amount of oxygen than the -ic acid is distinguished by the suf- 
fix -ous and its salts are called the -ites. For example, chlorous 
acid, HC102, and sodium chlorite. 

If there is an acid containing a still smaller amount of oxygen 
than the -ous acid, it is distinguished by adding to the name 
of the -ous acid the prefix .hypo- and its salts are called hypo- 
. . . . -ites; for example, hypochlorous acid and sodium hypo- 

The Preparation of the Oxygen Compounds of Chlorine. 
The first step in the preparation of the oxygen compounds of 
chlorine is the reaction between water and chlorine. 

H 2 + C1 2 <= HC1 + HC10 

From the hypochlorous acid so formed, hypochlorites can be 
prepared by neutralization with bases. From these hypo- 
chlorites by appropriate means chlorates, perchlorates, chlorites, 
and the oxides of chlorine may be made. The reaction between 
chlorine and water is therefore an important one and may 
appropriately be called the key reaction for the oxygen com- 
pounds of chlorine. It is reversible and is very incomplete 
under ordinary circumstances. Equilibrium results in this and 
other reversible reactions when the conditions described in what 
is known as the Law of Mass Action are fulfilled. In order 
that we may have a proper mastery of these reversible reactions, 
this law will now be developed. 

The Law of Mass Action. The law of mass action in the form 
in which it is to be developed here, applies to reversible reactions 
in equilibrium at constant temperature. Let us take a general 
case and say that we have a reversible reaction between A and B 
for the formation of C and D. Let a, b, c, and d /epresent the 
concentrations of A, B, C, and D respectively, and let R stand for 
the rate at which A and B combine for the formation of C and D, 


and B/ for that at which C and D react for the re-formation of A 
and B, 

A + B^C + D 


Then from what we have already learned, p. 44, the rate 
at which A and B react, everything else being constant, is pro- 
portional to a, and likewise to b, and therefore is proportional 
to their product since a number proportional to each of two 
or more numbers is proportional to their product or 

R = Ka 
R = K'b 
R = K"ab 

Correspondingly R' the rate at which C and D react is propor- 
tional to the concentration of C and to that of D and hence to the 
product of these concentrations 

R' = K'"c 
R' = K""d 
R' = K""'cd 

Equilibrium will result in all such reactions when the two rates 
are equal, and therefore at equilibrium 

R = R' 

K"ab = K'""cd 

K'"" ab " 

Expressed in words this would be: The product of the concentra- 
tion of the resulting substances divided by the product of the 
concentration of the reacting substances is a constant for a 
reversible reaction in equilibrium at constant temperature. From 
this it follows that if the three substances A, B, and C react to 
form D, E, and F, equilibrium will result when def/abc reaches a 
certain definite constant value whose magnitude depends upon 
the nature of* the substances and the temperature, but not upon 
the particular values of the concentrations which happen to exist 


in any one case. What is called the mass law equation for this 
reaction is: 


-T- = K 

Suppose now that A is identical with B, then "a" will be identical 
with "b." The chemical equation then becomes 

A + A + C<=D + E + F 

a a c d e f 

and the mass law equation, 

def = def = 
aac a 2 c 

If A is identical with B and C and if D and E are identical, these 
equations become, 

A + A + A<=D + D + F 

3A <= ^D + F 
ddf d 2 f 

o J\. 

This may be expressed in words as follows: When equilibrium 
results in a reversible reaction at constant temperature, the product of 
the concentration of the resulting substances divided by the product 
of the concentrations of the reacting substances each concentration 
raised to that power whose exponent is the coefficient of the substance 
in the chemical equation is a constant. The content of this 
statement is the Law of Mass Action. 

For an example of its application we may take the equilibrium 
between water vapor, iron, iron oxide, and hydrogen. The 
equation is 

3Fe + 4H 2 O <= Fe 3 O 4 + 4H 2 
a b c d 

The mass law equation is 

cd 4 

Since solids and liquids change their volume so slightly with 
changes of pressure, it follows that the ratio of their mass to their 


volume is practically constant; therefore the concentration of a 
solid substance or of a homogeneous liquid is constant and only 
gases and dissolved substances have variable concentrations. 

The iron and the iron oxide are solids so their concentrations 
are constant and may be combined with the equilibrium con- 
stant K. The conditions of equilibrium are therefore dependent 
upon the concentrations of the hydrogen and of the water vapor. 
The mass law equation then becomes 

cH Ka 3 
b 4 " : c 

This then tells us that these substances will be in equilibrium 
when there is a certain definite ratio between the concentrations 
of the hydrogen and of the water vapor. 

It is very easy to see from this that if the concentration of the 
hydrogen be decreased, the water must react with the iron for 
the formation of more hydrogen and iron oxide. In a current 
of steam which sweeps away the hydrogen, equilibrium will never 
be reached and the reaction will continue until all the iron is 
used up. 

The law of mass action is exceedingly useful in that it enables 
one to tell just what to do to make a reversible reaction run in 
either direction as desired, and further, if we know the magnitude 
of k, the value of the concentration of one substance can be 
calculated given that of the others. 

In accordance with the law of mobile equilibrium (page 57), 
chemical equilibrium is shifted in such a way that more and more 
of the substances, which are formed with the absorption of heat, 
will be present as the temperature rises. 

From this it may be seen that K, the equilibrium constant, 
varies with the temperature. At any one temperature it is, 
however, independent of the absolute value of the concentrations. 

Let us apply this mass law to the reaction between water and 
chlorine writing the latter ionically, 

H 2 o + ci 2 <= 2H+ + cr + cio- 

a b c d e 


In this a, b, etc., represent the concentrations of the substances 
just above them. Equilibrium will result when cMe/ab has 
reached a certain definite ratio or 

c 2 de _ 

From this it may be easily seen that a decrease in the concentra- 
tion of one of the substances on the right hand of the equation 
will decrease the numerator of the constant valued fraction, and 
this must result in the reaction proceeding from left to right, 
according to the above equation, until the values of a and b have 
been decreased and those of c, d and e have been increased suffi- 
ciently to restore the original equilibrium value for the fraction. 
By the addition of any hydroxide, say sodium hydroxide, the con- 
centration of the hydrogen as ion, "c," may be decreased through 
the formation of water, and practically all the free chlorine 
may be consumed. Carbonates, say sodium carbonate Na 2 CO 3 , 
or anything else which will use up the hydrogen ion will act in 
the same way. The carbonates use up the hydrogen ion through 
the formation of carbonic acid, H 2 CO3, which is not much more 
broken up into its ions than water. The H 2 CC>3 is unstable and 
breaks down into water and gaseous carbon dioxide which escapes 
from the liquid in bubbles. The equations are 

2H++Cl-+ClO-+2Na++2OH- = 

= H 2 0+C0 2 

The equations written in the ordinary way become 

HC1 + HC1O + 2NaOH = 2H 2 O + NaCl + NaCIO 

HC1 + HC1O + Na 2 CO 3 = H 2 + NaCl + NaCIO + CO 2 

In either case we say that sodium chloride and sodium hypo- 
chlorite are formed. The reaction is not usually carried out by 
saturating water with chlorine and then adding sodium hydroxide, 
NaOH, but by running chlorine into sodium hydroxide solution 
and the equation is often written, 

2NaOH + Cla = NaCl + NaClQ + H 2 O 


but the action doubtless takes place in the stages indicated 
above. Other hypochlorites can be made from the hydroxides 
of the corresponding metals in much the same way by the direct 
action of chlorine upon the hydroxide, either in solution or in the 
solid state. 

Without question the most important hypochlorite is that of 
calcium, which in the form of bleaching powder, or " chloride 
of lime," as it is often called, is manufactured by the thousands 
of tons from chlorine and calcium hydroxide. This bleaching 
powder is not a pure hypochlorite, but is a mixed salt, half 
chloride and half hypochlorite, formed by the union of one 
atomic weight of calcium with an atomic weight of chlorine ion 
and one of hypochlorite ion. The formula of the salt is given in 
various ways. Perhaps the most common of these are 

CaCl(ClO), Ca, or CaCl 2 O 

The equation for its formation is as follows, 

Ca(OH) 2 + C1 2 = CaCl(ClO) + H 2 O 

This bleaching powder is a solid and may be conveniently shipped. 
It is easily converted into chlorine, into sodium hypochlorite 
and chloride solution, or into hypochlorous acid. Upon passing 
into solution in water it acts as though it breaks up into calcium 
as ion and chlorine and hypochlorite ions. 

CaCl(C10)^Ca-H- + Cl~ + C1Q- 

Upon the addition of sodium carbonate the very slightly soluble 
calcium carbonate, CaC0 3 , is precipitated, and sodium, chlorine, 
and hypochlorite ions are left in solution. 

Ca+++Cl-+C10-H-2Na++C0 3 ^CaCO 3 +2Na++Cl-+C10- 


CaCl(ClO) + Na 2 CO 3 ^CaC0 3 + NaCl + NaCIO 

The solution, containing sodium hypochlorite, obtained in this 
way is very largely used for bleaching cotton and linen goods in 
the laundries. 

Solutions containing hypochlorite are often formed by elec- 


trolysis of solutions of chlorides. Chlorine is liberated at the 
anode and hydrogen with the simultaneous formation of hy- 
droxide at the cathode. The apparatus is so designed that the 
chlorine and hydroxide mix almost as soon as formed and hence 
hypochlorite is produced. 

The hypochlorites either alone or when mixed with chlorides 
are good oxidizing and bleaching agents and also are excellent 
disinfectants. One very important use for bleaching powder is 
in the purification of water. This depends upon the fact that 
very small quantities of the substance will kill typhoid and other 
pathogenic bacteria without injuring the water for domestic 

The ion C1O~ differs markedly in properties from the Cl~. It is 
for example, a strong oxidizing agent and does not give a precipi- 
tate of silver chloride upon the addition of silver nitrate. 

Decomposition of Hypochlorites. Hypochlorous acid and its 
salts are unstable and readily undergo change. Sodium hypo- 
chlorite, for example, slowly 'decomposes into the chloride and 
oxygen : 

2NaClO = 2NaCl + O 2 

The addition of a small quantity of a cobalt salt results in the 
formation of a black precipitate of cobaltic hydroxide. which acts 
as a vigorous catalyzer for the decomposition of the hypochlorites. 
In this way oxygen may be easily obtained from bleaching 
powder. By adding some cobalt salt to sodium hydroxide solu- 
tion and then passing in chlorine, a steady stream of oxygen may 
be obtained. This reaction may be represented as follows: 

4NaOH + 2C1 2 = 4NaCl + 2H 2 O + O 2 

Mixtures of iron and copper sulfates work in practically the 
same way as cobalt salts and are used in the technical preparation 
of oxygen from bleaching powder. 

This same decomposition of hypochlorites may be accelerated 
by light and takes place fairly rapidly in sunlight. The action of 
light on chlorine water finds an explanation from this fact. The 
chlorine acts on water, forming hypochlorous and hydrochloric 


acids, and the hypochlorous acid then decomposes into oxygen 
and hydrochloric acid: 

H 2 O + C1 2 <=HC1 + HC10 

2HC10 = 2HC1 + O 2 
or combining, 

2H 2 O + 2C1 2 = 4HC1 + 2 

This ability of hypochlorous acid and its salts to yield oxygen 
is closely connected with the oxidizing power of these compounds. 

In addition to decomposing into oxygen and chlorides, hypo- 
chlorites change slowly into chlorates and chlorides. For ex- 
ample, sodium hypochlorite passes into sodium chlorate and 
chloride, as shown in the following equation: 

3NaC10 = NaClO 3 + 2NaCl 

This change takes place more rapidly when the solution is heated 
and especially if the solution is acid or contains an excess of 
chlorine. An alkaline solution of a hypochlorite is fairly stable. 
The chlorate formed as described above slowly changes into 
chloride and perchlorate and the latter in turn decomposes into 
chloride and oxygen. In order that the reactions shall take 
place at a reasonable rate the temperature must be far above 
the boiling-point of water, so to obtain them the dry salts are 
heated. The equations are 

* 4NaClO 3 = 3NaC10 4 + Nad 

NaC10 4 = Nad + 20 2 

The final result is that, after going through these stages, the 
hypochlorite ultimately is changed completely into chloride and 
oxygen as is the case with the cold solution in the presence of a 
cobalt salt. The hypochlorite, chlorate, and perchlorate are 
unstable substances, while the chloride and oxygen are stable. 
Except in the presence of cobalt oxide, etc., the system sodium 
hydroxide and chlorine does not pass directly to its most stable 
state, sodium chloride, oxygen and water, but goes through a 
series of stages each more stable than its predecessor until the 
final stable condition is reached. In this respect, this system 
seems to be typical of many others which behave similarily. So 



generally is this the case that it is described in what is called 
Ostwald's Law of Successive Reactions, which states that a system 
does not pass directly from the least stable to the most stable state, 
but does so through a series of steps or stages of gradually increasing 

We may represent the successive steps in the transformation 
of the unstable system, sodium hydroxide and chlorine, into the 
stable one, sodium chloride, oxygen and water as follows: 


12H 2 O + 12NaCl + |l2NaClO| 


NaCl + |3NaC10 


12NaCl + 6O 2 4NaCl + 6O 2 3NaCl + 6O 2 

The conditions stated along each arrow show the factors which 
favor the change indicated by the arrow. 

This is typical of the changes undergone by other systems of 
bases and chlorine and gives briefly a general scheme for the 
preparation of hypochlorites, chlorates, and perchlorates. 

Hypochlorous Acid. Hypochlorous acid is very volatile and 
will stand distillation, consequently it is easily prepared from 
its salts by the first general method, by adding to a hypochlorite 
a less volatile acid and distilling. The reaction must be carried 
out in rather dilute solution since the concentrated acid readily 
decomposes into water and chlorine monoxide, QUO. Since 
hypochlorftes are difficult to prepare pure, a mixture of hypo- 
chlorite and chloride is generally used. Hypochlorous acid 
HC10, is a very much weaker acid than hydrochloric, HC1 
and is almost exclusively formed if a little less acid is added than 
the amount required to act upon the hypochlorite which is 
present. It is very important that an excess of acid be avoided 
since hydrochloric and hypochlorous acids react very rapidly to 
form chlorine and water. 

HC1 + HC10 = H 2 + C1 2 


Bleaching powder is usually used to furnish the hypochlorite 
and nitric acid is the added acid, care being taken to add it a 
little at a time and to stir vigorously to avoid any local excess. 
The equation is 

2CaCl(C10) + 2HNO 3 = 2HC10 + CaCl 2 + Ca(NO 3 ) 2 

Instead of nitric acid one may use hydrochloric or still better, 
the very weak boric acid. 

The hypochlorous acid formed is volatile and passes over with 
the first part of the water. It may be still further concentrated 
by fractional distillation. 

Properties of Hypochlorous Acid. The anhydrous acid cannot 
be prepared, so it is known only in solution. Its concentrated 
solutions have a yellow color and smell something like bleaching 
powder. In dilute solutions it is almost colorless. It neutralizes 
strong bases and forms hypochlorites. 

It is unstable and decomposes in several ways. When warmed 
or exposed to sunlight, it may decompose into hydrochloric acid 
and oxygen. 

2HC1O = 2HC1 + O 2 

or it may change into hydrochloric and chloric acids. 
3HC10 = 2HC1 + HC10 3 

This latter change takes place even in the dark. 

In the two cases given above, hypochlorous acid acts like other 
hypochlorites, but hypochlorous acid differs in one respect from 
other hypochlorites, and that is, in very concentrated solution it 
will decompose into water and chlorine monoxide. 

2HC10<=H 2 O + C1 2 

This reaction is reversible and hypochlorous acid is formed by 
the union of water and chlorine monoxide; consequently the 
latter is often called hypochlorous anhydride. An anhydride 
is an oxide which will combine with water to form an acid. 

As has been mentioned, hypochlorous acid is a powerful 
oxidizing and bleaching agent, being more active in this respect 
than other hypochlorites. In fact, there is some little evidence 
that solutions of hypochlorites owe a part of their bleaching power 


to the presence in them of hypochlorous acid. Hypochlorous 
acid bleaches because it oxidizes and destroys the coloring matters 
which are usually very complex carbon compounds. Each of 
these compounds owes its color to a certain definite combination 
of its constituents, and when this combination is broken up by 
oxidation or otherwise the color is destroyed. Usually the 
reaction is too complex to be represented by an equation. 

It will be recalled that chlorine was not active as a bleaching 
agent until the article to be acted upon was moistened. It is 
probable that it is really the hypochlorous acid formed by the 
action of the chlorine on the water which does the bleaching. 
Cotton or linen goods are prepared for bleaching by boiling them 
in a soap solution for a time to remove grease and oils and then 
passing them through a dilute solution of bleaching powder 
and into very dilute sulfuric acid. The last two processes are 
repeated until the desired results are reached. The goods are 
then carefully washed to remove the excess of chemicals. The 
solutions must be very dilute and the operation carried out with 
care or a large part of the strength of the fabric will be destroyed. 
Even under the best conditions the cloth is considerably weak- 
ened. Wool, silk, and feathers are so greatly altered by chlorine, 
hypochlorous acid and hypochlorites, that they cannot be 
bleached by these substances. Hypochlorous acid is formed by 
the action of carbon dioxide, C02, upon moist bleaching powder, 

2CaCl(C10) + H 2 + C0 2 = CaC0 3 + CaCl 2 + 2HC1O 
and sometimes in bleaching, the goods which have been wet in 
bleaching powder solution, are simply hung in the air instead of 
being treated with sulphuric acid. Hypochlorous acid is- a power- 
ful disinfectant. 

Chlorine Monoxide. Chlorine monoxide or hypochlorous 
anhydride is usually prepared not by the decomposition of 
hypochlorous acid which was mentioned above, but by the 
action of chlorine on mercuric oxide whereby mercuric oxy- 
chloride and the chlorine monoxide are formed according to the 
following equation: 

2HgO + 2C1 2 = C1 2 + HgOHgCl 2 

The chlorine monoxide is a yellowish-brown gaseous substance 
which may be condensed to a liquid boiling at + 5C. Its odor 


is strong but distinctly different from that of chlorine. Both 
the gas and the liquid are explosive. The gaseous substance is 
easily soluble in water. The solution contains hypochlorous 
acid and therefore this monoxide is also called hypochlorous 

Chlorates.-^Chlorates are easily formed by heating solutions 
of hypochlorites as has been described above. The reactions 
take place most readily when the solutions are concentrated 
and slightly acid or contain an excess of chlorine. 

Potassium chlorate is the most widely used and important of 
these salts, although sodium chlorate is taking its place for 
many purposes because the latter is more soluble. 

Potassium chlorate may be prepared by passing an excess of 
chlorine into a warm concentrated solution of potassium hydrox- 
ide. The reaction takes place in stages as in the formation of 
sodium chlorate, the hypochlorite being first formed. The 
equations are 

2KOH + C1 2 = KC1 + KC10 + H 2 O 
3KC10 = 2KC1 + KC10 3 

or combining into one and omitting the intermediate stage, 
6KOH + 3C1 2 = 5KC1 + KC10 3 + 3H 2 O 

The chlorate is much less soluble than the chloride, so by con- 
centrating the solution and allowing it to cool, the chlorate 
separates out mixed with only a little of the chloride, which may 
be removed by a second crystallization. As will be seen from the 
equation just given, only one-sixth of the somewhat expensive 
potassium hydroxide is converted into chlorate. The chloride 
formed at the same time, occurs in nature in large quantities and 
has a comparatively small value. On a manufacturing scale, this 
waste is avoided by first preparing calcium chlorate, Ca(C103) 2 , 
by the action of an excess of chlorine upon hot calcium hydroxide. 

6Ca(OH) 2 + 6C1 2 = 5CaCl 2 + Ca(C10 3 ) 2 + 6H 2 O 

Potassium chloride is then added to the solution in a little more 
than the required quantity to react with the calcium chlorate 
according to the following equation. 

Ca(C10 3 ) 2 + 2K01 <F CaCl 2 + 2KClO a 


This reaction is reversible and far from complete, but the potas- 
sium chlorate is only moderately soluble in cold water, while the 
potassium chloride is easily and the calcium chloride and chlorate 
exceedingly soluble; so when the solution is concentrated and 
cooled, potassium chlorate passes out of solution in the form of 
crystals. As the potassium chlorate separates from the solution 
more is formed until finally the greater part of the chlorate has 
been obtained in the form of the potassium salt. 

Chlorates are also very largely made by electrolysis, the process 
being very much like that for the preparation of hypochlorites, 
the main difference being that the solution is kept hot and very 
slightly acid, conditions which it will be recalled are favorable 
to the change of hypochlorites into chlorates. 

Properties of Chlorates. All chlorates are at least moderately 
soluble in water, and many are so very soluble that they are 
deliquescent. Potassium chlorate is one of the least soluble of 
these salts, 100 parts of water at 0C. dissolve 3.14 parts of this 
salt. The solubility increases rapidly with rising temperature so 
that at 100C. it is approximately eighteen times that at 0C. 

The chlorates all give off oxygen upon being heated and are 
strong oxidizing agents. Because of this property they are 
much used in making fireworks, dyes, matches, and explosives. 
An intimate mixture of sugar, C^H^Ou, and potassium chlorate 
is sometimes used as a blasting powder, but like practically all 
other explosives containing chlorates, it is treacherous and 
dangerous to handle. The addition of a little castor oil makes 
these explosives safer. 

Potassium chlorate, it will be recalled, was used in the labora- 
tory preparation of oxygen. The equation is, 

2KC10 3 = 2KC1 + 30 2 

The oxygen so prepared almost always contains a trace of 

Like sodium chlorate, potassium chlorate if slowly and care- 
fully heated will pass into potassium perchlorate, KC10 4 : 

4KC10 3 = KC1 + 3KC10 4 

Although chlorates are good oxidizing agents, they are less 
powerful than the hypochlorites being more stable compounds. 


Preparation of Chloric Acid. Chloric acid cannot be distilled 
without decomposition and therefore cannot be prepared from 
its salts by the first general method for the preparation of acids. 
It is made by what is known as the second general method for the 
preparation of acids from their salts which is briefly as follows: 
Add to a solution of a salt of the desired acid another acid which is 
so chosen that the anion of the added acid shall form a very diffi- 
cultly soluble compound with the cation of the salt. This com- 
pound is at once formed and passes out of solution, leaving the 
hydrogen ion of the added acid and the anion of the salt in solu- 
tion, that is, forming a solution of the desired acid. By careful 
evaporation, this acid may be then obtained in a pure or at least a 
more or less concentrated form. Of course the salt to be used 
and the acid to be added must be very carefully chosen, and just 
the correct amount of each used. Silver chloride and barium 
sulfate are among the least soluble of the salts, so silver chlorate, 
AgClO 3 , with hydrochloric acid, or barium chlorate, Ba(C10 3 ) 2 , 
with sulfuric acid are well adapted for the formation of chloric 
acid according to the following equations: 

AgC10 3 + HC1 = HC10 3 + AgCl 

Ba(ClO 3 ) 2 + H 2 SO 4 = 2HC1O 3 + BaSO 4 

Barium chlorate is much cheaper than silver chlorate, so the 
latter reaction is the one usually used. The acid solution is 
filtered from the white precipitate of barium sulfate, BaSO 4 , and 
concentrated by evaporation at a temperature below 40C. 
until it contains 40 per cent, of the acid. The dilute solution is 
colorless while the concentrated has a yellowish tint. It is a 
strong acid. 

The solution slowly decomposes in a number of ways. It may 
form perchloric acid, HC104, 

4HC10 3 = HC1 + 3HC1O 4 

or break down into hydrochloric acid and oxygen, 
2HC10 3 = 2HC1 + 3O 2 

The HC1 so formed then reacts with more HC10 3 to form chlorine 
and water, 

5HC1 + HC10 3 = 3H 2 + 3C1 2 


In addition it may go over into water, chlorine dioxide, C1O 2 , and 

4HC10 3 = 2H 2 O + 4C1O 2 + 2 

When concentrated sulfuric acid is poured on a chlorate, chloric 
acid is formed, which at once decomposes as shown above. The 
chlorine dioxide so produced generally explodes spontaneously 
with such violence that only very small quantities of chlorate 
should be used. 

Chloric acid has found no important application, but its salts 
are largely used. 

Separation of Salts. A solid salt generally separates from its 
solutions in the form of crystals, and the process of obtaining 
these is called crystallization. It is very evident from what has 
been learned, p. 61, that crystals cannot form in an unsaturated 
solution-. In fact the solution must be more than saturated, 
supersaturated, before they can separate. But in the absence 
of the solid solute a solution may be considerably supersaturated 
without crystallization taking place. 

Since most substances are more soluble at high temperatures 
than at low, it is usually very easy to prepare supersaturated 
solutions by making a saturated solution at a somewhat elevated 
temperature and then cooling the liquid after the removal of 
any excess of the undissolved solid. The cooled solution becomes 
supersaturated and may often be preserved for a long time in this 
condition by carefully excluding dust which usually contains 
enough small particles of the solute to bring about crystallization. 
Generally, however, it will spontaneously deposit crystals as 
soon as it reaches a certain degree of supersaturation. After the 
separation has once started it will continue until only so much of 
the solute remains dissolved as corresponds to its solubility at 
that temperature. If the supersaturated solution was prepared 
from crystals of the substance, the process of reobtaining the 
crystals is known as recrystallization, and is very commonly used 
in the purification of substances. The process of recrystallization 
which has just been outlined is especially adapted to the purifica- 
tion of such substances as change their solubility very rapidly 
with the temperature. Potassium chlorate is a good example. 
The process is not so well adapted to potassium chloride, and is 



of practically no value for things like sodium chloride which are 
almost as soluble in cold water as in hot (see Fig. 26). In such 
cases the only way to recrystallize the substances is to evaporate 
a part of the solvent which will of course cause a corresponding 
amount of the solute to crystallize. Obviously all of the solvent 
should not be evaporated since this would leave the salt as 
impure as before. 

When a mixture of salts is dissolved in water,' the numerical 
value of the solubility of each salt is more or less altered but the 
general relationships are left unchanged. Before any given 
salt can crystallize, the solution must become supersaturated 
with that particular salt. As a general rule each salt tends to 


I 40 






20 c 

30 c 

40 50 60 C 

FIG. 26. 


90 100 

crystallize by itself, and if two or more separate at the same 
time, a mixture of crystals of the salts will be obtained rather 
than crystals, each of which contains every salt. 

Upon the evaporation of a solution of a mixture of two salts, 
say for example potassium chloride and chlorate, that salt will 
separate first whose point of saturation is first exceeded; in this 
case it is the chlorate at ordinary temperatures, if the salts are 
present in anything like equal quantities. By removing the 
crystals as they are formed, they may be obtained in an approxi- 
mately pure state. Ultimately as the concentration proceeds, 
the solution will become saturated with the other salt, potassium 
chloride, and a mixture of crystals will then be formed. 


It is not necessarily the least soluble salt which separates first, 
but that one whose point of saturation is first exceeded. This 
may easily be the case with the more soluble salt if it is present 
in relatively large quantities. 

Perchlorates. Potassium perchlorate is easily prepared by 
heating potassium chlorate until it melts and then keeping it at 
this temperature for some time. A little oxygen is given off and 
solids begin to form in the molten salt. After a time the 
whole thing goes solid, although the temperature has not been 
lowered. The solid so formed is a mixture of potassium chloride, 
KC1, and potassium perchlorate, KC10 4 . It is formed according 
to the following equation, 

4KC1O 3 = 3KC1O 4 + KC1 

The evolution of the oxygen had nothing essentially to do with 
this reaction, and hence is not shown in the equation. It is 
formed by the side reaction, 

2KC1O 3 = 2KC1 + 30 2 

The potassium perchlorate, KC1O 4 , is only slightly soluble, so 
it is easily separated from the potassium chloride, KC1, by finely 
powdering the mixture and dissolving out the potassium chloride 
with cold water. A recrystallization of the residue from hot 
water yields pure potassium perchlorate, KC104. 

The perchlorates are much more stable than the hypochlorites 
or chlorates. Upon being heated they yield oxygen and chlorides, 
and may also act as oxidizing agents but are not as active as 
hypochlorites or chlorates. They are all more or less soluble in 
water, the potassium salt being one of the least soluble. Some 
of them are used in matches and fireworks because they are safer 
than the chlorates. 

Perchloric Acid. Perchloric acid is easily formed from the 
potassium salt by the first general method, using sulfuric acid 
and cautiously distilling under diminished pressure. Dense 
white fumes are evolved and a colorless or slightly yellow liquid 
distills, consisting of pure perchloric acid, HC1O 4 , which condenses 
in the receiver. 

Mixtures of perchloric acid and water act like mixtures of 
hydrochloric acid and water in that there is a constant boiling 


mixture of minimum vapor pressure and maximum boiling-point. 
This mixture contains 72 per .cent, of acid and boils at 203C. 
By adding to this constant boiling mixture twice its volume of 
concentrated sulfuric acid, and distilling, pure perchloric acid 
may be obtained. 

The pure acid is volatile, boiling at 19C. under a pressure of 
1.1 cm. of mercury; and has a density of 1.78. It is a good 
oxidizing agent and decomposes with great readiness so that 
when brought in contact with oxidizable substances it generally 
explodes with violence. When heated under atmospheric pres- 
sure it decomposes at 92C. It dissolves in water with a hissing 
noise and the evolution of much heat. The aqueous acid of all 
strengths up to the constant boiling mixture is entirely stable. 
Unlike hypochlorous and chloric acids, it does not liberate 
chlorine from hydrochloric acid. It combines with water to 
form white crystals of perchloric acid monohydrate HC1O 4 -H 2 O. 

Perchloric Anhydride. Perchloric anhydride, C1 2 O 7 , may be 
prepared by adding to perchloric acid phosphorus pentoxide, 
P 2 O5, keeping the whole cool by a freezing mixture of ice and 
salt. After standing for a day the mixture is distilled. The 
anhydride is colorless and boils at 82C. It may be distilled 
under ordinary pressure without danger, but explodes when 
struck or brought in contact with a flame. 

Chlorine Dioxide and Chlorites. As has already been men- 
tioned, one of the ways in which chloric acid decomposes, is into 
water, chlorine dioxide, C1O2, and oxygen: 

4HC1O 3 = 2H 2 O + 4C1O 2 + O 2 

Chlorine dioxide is a yellow-brown gaseous substance which 
may be condensed in a tube cooled with ice and salt. The boil- 
ing-point of the liquid is about 10C. Both the liquid and its 
vapor are extremely explosive. The vapor explodes when 
brought in contact with a metal rod heated to a temperature 
far below redness. The products of the explosion are chlorine 
and oxygen. 

When passed into sodium hydroxide solution, it yields a mixture 
of sodium chlorate and chlorite, NaC10 2 . 

2NaOH + 2C1O 2 = NaC10 3 + NaClO 2 -j- H 2 O 


The chlorites are all unstable substances for which no appli- 
cations have been found. The acid is known only in dilute solu- 
tions. The chlorites appear to stand between the hypochlorites 
and the chlorates in the order of stability, but the evidence is not 
entirely conclusive. 

Relative Oxidizing Power. It might very reasonably be 
thought that perchloric acids and chloric acids would be 
stronger oxidizing agents than hypochlorous acid, because the 
former contain more oxygen per mole; but this is not the case. 
Hypochlorous acid is the strongest, and perchloric the weakest. 
This is connected with the fact that hypochlorous acid is the 
most unstable and perchloric the most stable of the series. 


General. Fluorine, chlorine, bromine and iodine form a group 
or family of elements which closely resemble one another in their 
chemical properties and those of the compounds which they form. 
This group is called the halogen family. From what is to follow, 
it will be seen that there is a regular change in properties of the 
halogens and their compounds as the atomic weight of the 
element increases. It will also be very evident that chlorine, 
bromine and iodine resemble each other much more closely than 
they do fluorine. For this reason the discussion of chlorine will 
be followed immediately by that of bromine and iodine and 
fluorine will be taken up last. / 

Bromine and iodine are so much like chlorine that a very fair 
knowledge of their chemistry may be easily obtained by compari- 
son with that of chlorine, noting the many points of resemblance 
and the few differences. The root of most of these differences lies 
in the fact that bromine is liberated from its compounds by weaker 
oxidizing agents than chlorine, and that iodine is still more readily 
set free than bromine. 


Occurrence. Like chlorine, bromine is never found free, but 
only in combination, chiefly as bromides of sodium, potassium, 
and magnesium. It is not nearly as abundant as chlorine, and 
its compounds are never found in anything like a pure state, but 
are always mixed with very much larger quantities of chlorides. 
It is found in sea water, many mineral springs, salt wells, and beds 
of rock salt. These beds of rock salt have doubtless been formed 
by the evaporation of sea water, so the presence of bromides in 
them is not surprising. The bromine content of salt beds differs 
rather widely, and naturally only the richer deposits are utilized 
as sources of bromine. The chief bromine-producing localities 




are Stassfurt in Germany, and Michigan, Pennsylvania, and 
West Virginia, in the United States. 

Preparation of the Element. From the fact that bromine is 
more easily liberated than chlorine, it follows that any of the 
oxidizing agents used in the preparation of chlorine may be 
employed for bromine, so we may match each of the methods 
given in the preceding chapter for the preparation of chlorine by 
a similar one for bromine, as shown by the following equations : 

4HC1 + O 2 = 2H 2 O 
4HBr -f 2 = 2H 2 O 
4HC1 + MnO 2 = MnCl 2 
4HBr + MnO 2 = MnBr 2 

2C1 2 
2Br 2 
C1 2 
Br 2 

2H 2 O 
2H 2 O 

2KMnO 4 
CaCl 2 O 
CaCl 2 

2H 2 
2H 2 O 
8H 2 
8H 2 O 

2NaCl + Mn0 2 + 3H 2 SO 4 = MnS0 4 + 2NaHS0 4 + C1 2 
2NaBr + Mn0 2 + 3H 2 S0 4 = MnS0 4 + 2NaHS0 4 + Br 2 
2KMn0 4 + 16HC1 = 2MnCl 2 + 2KC1 + 5C1 2 
16HBr = 2MnBr 2 + 2KBr + 5Br 2 
2HC1 = CaCl 2 + C1 2 + H 2 
2HBr = CaCl 2 + Br 2 + H 2 

Electrolysis of chlorides and bromides yields respectively 
chlorine and bromine at the 

Of these reactions, that be- 
tween the bromides, manganese 
dioxide and sulfuric acid (Fig. 
27), is the most important be- 
cause the bromides are much 
cheaper and more easily ob- 
tained than hydrobromic acid. 

In addition to the above, 
bromine is liberated by a num- 
ber of oxidizing agents which 
cannot act on chlorides. The 

element chlorine is an oxidizing agent, not as powerful of course 
as the ones used to set it free, but still powerful enough to 
liberate bromine as shown in the following equation: 

2NaBr + C1 2 = 2NaCl + Br 2 

Sodium Bromide, 1=5 
Manganese Dioxide, ^~^f : 
Sulfuric Add Bromine 

FlG 2 7. 


The reaction is really between the free chlorine and the bromine 

2Br~ .+ C1 2 = 2C1- + Br 2 

and consequently is an exchange of charges, and yet the liberation 
of bromine from a bromide is considered oxidation; so this kind 
of exchange of charges must be counted oxidation and kept in 
mind until a general definition of oxidation and reduction can be 
given. It is by this reaction that most of the bromine of com- 
merce is obtained. 

Concentrated sulfuric acid is a fairly strong oxidizing agent, 
strong enough to act on hydrobromic acid, but not on hydro- 
chloric acid. 

2HBr + H 2 SO 4 = Br 2 + S0 2 + 2H 2 O 

This last action must be kept in mind in the preparation of 
hydrobromic acid. 

Physical Properties of Bromines-Bromine is a dark, brownish- 
red liquid, which is almost opaque in thick layers, and has a 
density of 3.19. It boils at 59C. and freezes at - 7.3C. The 
liquid produces very severe "burns" if it comes in contact with 
the skin. Bromine is an easily volatile liquid since its vapor 
pressure at ordinary temperature is about 20 cm. The vapor 
has a deep red color, an exceedingly stifling odor, and is very 
irritating to the mucous membrane. The name bromine is de- 
rived from the Greek word, /3pco/xos, meaning a stench. The den- 
sity of the bromine vapors is five times that of oxygen and in- 
dicates a molecular weight of 160. Since the atomic weight of 
bromine (79.92) is half the molecular weight, its formula is Br 2 . 
At higher temperatures the molecular weight becomes somewhat 
less, indicating a partial dissociation of the Br 2 into Br. Iodine 
shows this tendency still more strongly. 

Bromine dissolves in water to the extent of about 3 parts by 
weight of bromine to 100 parts of water at ordinary temperature. 
The solution has a yellowish-red color, smells of bromine, and 
may often be used in the place of the latter when water does 
not interfere with the reactions. The solubility of bromine 
is markedly increased if the water also contains some soluble 
bromides. It is more soluble in carbon disulfide, alcohol, and 
many organic solvents than it is in water. 


Chemical Properties. The chemical properties of bromine are 
very much like those of chlorine, modified by the fact that it is 
slower and less intense in its action than the latter. It is a 
powerful oxidizing agent, though less so than chlorine. It 
combines directly with hydrogen and most of the other elements 
forming bromides. With water it forms a hydrate, Br 2 -8H 2 0, 
which, in general, is similar to chlorine hydrate. 

Hydrogen Bromide. With slight modifications due to the facts 
mentioned above, that bromine does not combine as vigorously 
with hydrogen as chlorine does, and that hydrogen bromide is 
more easily oxidized than hydrogen chloride, practically every 
method given in Chapter VIII for the preparation of hydrogen 
chloride may be matched by a corresponding one for hydrogen 
bromide. This is shown by the following equations. 

H 2 + C1 2 = 2HC1 
H 2 + Br 2 = 2HBr 

The mixture of hydrogen and bromine does not explode when 
exposed to sunlight or an electric spark, as that of hydrogen and 
chlorine does, but combines fairly rapidly at high temperatures 
under the influence of platinum as a catalyzer. 

Nad + H 2 SO 4 = HC1 + NaHSO 4 
NaBr + H 2 SO 4 = HBr + NaHSO 4 

In this case dilute sulf uric acid must be used in the preparation of 
hydrogen bromide because the concentrated acid is a strong 
enough oxidizing agent to liberate the bromine. 

The indirect preparation of hydrogen bromide by the action of 
bromine upon compounds of hydrogen and carbon corresponds 
closely to that of hydrogen chloride by the action of chlorine 
upon these same compounds. For example, bromine and ben- 
zene, C 6 H 6 , react for the formation of hydrogen bromide and 
monobrombenzene, C 6 H 5 Br. 

Br 2 + C 6 H 6 = HBr + C 6 H 5 Br 

The method which is generally the most convenient for the 
preparation of hydrogen bromide is by the interaction, at ordi- 
nary temperatures, of bromine, red phosphorus and water (Fig. 28). 
The action takes place in steps. First the bromine and the 



phosphorus combine for the formation of phosphorus tribromide, 
PBr 3 . 

2P + 3Br 2 = 2PBr 3 

This then reacts with water to form phosphorous acid, H 3 PO 3 , 
and hydrogen bromide, HBr. 

PBr 3 + 3H 2 = 3HBr + H 3 PO 3 

Hydrogen chloride might be prepared in an analogous way, but 
the action is very violent and the method is not so cheap and 
convenient as that from salt and sulfuric acid. 

''Purl f ing Tube 
containing /%?/?/ 
Red Phosphorus 

FIG. 28. 

Properties of Hydrogen Bromide. Like hydrogen chloride, 
hydrogen bromide is a colorless gaseous substance at ordinary 
temperatures. Its critical temperature is 91C. At low tem- 
peratures it condenses to a colorless liquid, boiling at 68.5C. 
and freezing at 86C. The gaseous density of hydrogen bro- 
mide shows that it has a molar weight of 81, and the formula 
HBr. It is very soluble in water, 600 volumes of hydrogen 
bromide to one of water, and fumes upon contact with moist air. 
Like hydrogen chloride, it forms a maximum boiling solution 
with water. The composition of this is 48 per cent, hydrogen 
bromide, and it boils at 126C. 

Like hydrogen chloride, dry hydrogen bromide is a rather 
indifferent substance, but when in solution in water it is as 


highly acid as hydrochloric acid and its solution is known as 
hydrobromic acid. It acts upon metals, much as other acids do, 
forming hydrogen gas and solutions of the bromides of the metals. 
Speaking in terms of the ionic theory, solutions of hydrobromic 
acid and bromides contain the bromine^ ion, Br~. The properties 
of this ion are very much like those of the chlorine ion. It is 
colorless, forms a very difficultly soluble compound with silver, 
silver bromide, AgBr, which is less soluble than silver chloride 
and has a slight yellow color. The liberation of bromine from 
the bromine ion by chlorine is the easiest way to distinguish 
the bromine ion from the chlorine ion. The equation for the 
reaction is 

2H+ + 2Br~ + C1 2 = 2H+ + 2C1- + Br 2 

The free bromine colors the solution yellowish-red. If the solu- 
tion be shaken with a little carbon bisulfide, the latter takes up 
most of the bromine from the water and becomes highly colored. 
This makes a very delicate test for bromine. 

Oxygen Compounds of Bromine. No one has, so far, succeeded 
in preparing oxides of bromine, perbromic acid, or perbromates, 
but hypobromous acid, HBrO, and bromic acid, HBrO 3 , and 
their salts are well known. The reactions for their preparation 
are exactly analogous to those for the corresponding chlorine 
compounds. The first step is the reversible reaction between 
bromine and water for the formation of hydrobromic and hypo- 
bromous acids, 

Br 2 + H 2 O^HBr + HBrO 

When equilibrium results, the concentrations of the acids are 
even less than in the corresponding reaction for chlorine. The 
addition of sodium or potassium hydroxides will destroy the 
equilibrium by neutralizing the acids and cause the reaction to 
run toward the right until nearly all the bromine has been used 
up, and there results a solution of a mixture of the bromide and 
hypobromite salts. 

HBr + HBrO + 2NaOH = NaBr + NaBrO + 2H 2 O 
or, combining the two equations 

Br 2 + 2NaOH = NaBr + NaBrO + H 2 O 


The change of the hypobromites into bromates takes place 
even more rapidly upon heating, than the corresponding change 
of the hypochlorites. The equation is 

SNaBrO = 2NaBr + NaBr0 3 

Hypobromous acid is so unstable that even a dilute solution of 
it is hard to make and soon decomposes. Bromic acid may be 
made from its salts by the second general method for the prepara- 
tion of acids, or by the action of hypochlorous acid upon bromine 
water : 

10HC1O + Br 2 = 5C1 2 + 2HBrO 3 + 4H 2 O 


Iodine occurs in nature almost exclusively in combination, 
chiefly as iodides and iodates. It has, however, been found free 
in waters of certain mineral springs, having been liberated from 
hydriodic acid by the action of dissolved oxygen. This is in line 
with the fact that iodine is more easily liberated by oxidizing 
agents than either chlorine or bromine. 

Iodides are found in sea water, but in much smaller quantities 
than bromides. Fairly large quantities of iodates, chiefly that 
of sodium, are found in the Chile saltpeter beds, and the greater 
part of the iodine of commerce is obtained from the mother liquor 
left in the purification of this crude sodium nitrate. The other 
source of iodine is from ashes of certain seaweeds gathered off 
the coasts of Scotland and Ireland. They were formerly the main 
source of iodine, and it was in their ashes that the element was 

Preparation of the Element. Iodine may be liberated from 
iodides by any of the oxidizing agents used for chlorine or 
bromine, and by some others in addition which are too feeble to 
serve in the other cases. A few of the equations are given 
below; others may be easily written by simply substituting the 
symbol "I" for "Cl" or "Br" in the corresponding equations for 
these elements: 

4HI + O 2 = 2H 2 + 2I 2 
4HI + Mn0 2 = MnI 2 + I 2 + 2H 2 O 
2KI + Mn0 2 + 3H 2 S0 4 = MnS0 4 + 2KHSO 4 + Ij + 2H 2 

The last reaction is very commonly used (Fig. 29). 



"of Iodine 

Like bromine, iodine may be liberated by the action of chlorine 
or strong sulfuric acid upon hydriodic acid or iodides: 

2KI + C1 2 = 2KC1 + I 2 
SHI + H 2 SO 4 = 4H 2 + H 2 S + 4I 2 

Hydriodic acid is oxidized even more readily by sulfuric acid 
than hydrobromic is, and hence it is not possible to prepare 
hydriodic acid by the distillation of an 
iodide with sulfuric acid. 

Just as free chlorine will liberate free 
bromine from hydrobromic acid or bro- 
mides, so bromine will set iodine free from 
the corresponding iodine compounds: 

2KI + Br 2 = 2KBr + I 2 

Physical Properties of Iodine. At 

ordinary temperatures chlorine is gaseous, 

bromine liquid, and iodine solid; this shows 

a regular change in the physical state 

as the atomic weight increases. Iodine 

crystallizes in large rhombic plates, having 

a steel gray color and an almost metallic 

luster. It melts at 114C., and boils at 

184C. The vapor has a beautiful violet color which is so 

strong that it can be seen in a bottle containing iodine at 

ordinary temperatures, although the vapor pressure of iodine at 

such temperatures is very low. 

The vapor density of iodine at ordinary pressures and tempera- 
tures below 700C. shows a molecular weight of 254. The atomic 
weight of iodine is 126.92, hence the formula for iodine under 
these conditions is I 2 . If the temperature be raised above 700C. 
the molecular weight gradually decreases until at 1,700C. it 
becomes constant again at 127 corresponding to the formula I. 
Rise in temperature then above 700C. has caused a reversible 
dissociation of I 2 into I. 

FIG. 29. 

The color of I is different from that of I 2 , and doubtless if the 
high temperature did not interfere with its study I would be 
found to be a strikingly different substance from I 2 . 


Iodine is but very slightly soluble in pure water, imparting to 
it a brownish color. It is much more soluble in alcohol, chloro- 
form, and carbon disulfide than in water. The alcoholic solution 
is brown, while that in chloroform or carbon disulfide has the 
violet color of the vapor. Free iodine will dissolve in solid starch 
forming a solid solution having a strong blue color. This may 
be used as a very delicate test for either starch or iodine. In the 
latter case, it is carried out as follows: a suspension of finely 
divided starch in water is prepared by pouring boiling water over 
starch rubbed to a cream with a little water. Some of this is 
added to the solution suspected to contain iodine. If the latter 
is free, the strong blue color of the starch iodide solution at once 
appears. If the iodine is not free, it is liberated by the cautious 
addition of a little chlorine water. 

Law of Distribution. As mentioned above, iodine is much 
more soluble in carbon disulfide than in water. Carbon disulfide 
and water are not appreciably soluble in each other, and when 
brought together form two liquid layers, the carbon disulfide at 
the bottom. If now, some iodine be introduced, it will dissolve 
both in the water and the carbon disulfide, but to a far larger 
concentration in the latter. In fact, when equilibrium results it 
will be found that the weight of iodine per cubic centimeter in 
the carbon disulfide is 600 times that per cubic centimeter in the 

This is an example of the phenomena described by the law of 
distribution which states that: When two immiscible liquids are 
brought in contact with a third substance, soluble in each, this third 
substance distributes itself between the two solvtnts until its concen- 
tration in the one solvent bears a certain constant ratio to its concen- 
tration in the other. In this case, the ratio of the concentration 
in the carbon disulfide to that in water is 600 : 1. Some sub- 
stances have higher molecular weights in certain solvents than in 
others and this law holds only when the molecular weight of the 
third substance is the same in each of the solvents. 

Chemical Properties of Iodine. The chemical properties of 
iodine are very much like those of chlorine and bromine except 
that it is much less active and is not so strong an oxidizing agent. 
It combines directly with hydrogen and many other elements for 
the formation of iodides. It does not form a hydrate with water. 


Iodine is much more soluble in solutions of iodides than in pure 
water. In many ways the solution acts as though the iodine 
were free, in others, it behaves as if it were in combination. 
This is explained by assuming the existence in the solution, of a 
complex ion, I 3 ~ formed by the reversible reaction shown below: 

r + I 2 <= la- 
Hydrogen Iodide. The slight differences in the methods for the 
preparation of hydrogen iodide from those for hydrogen chloride 
and bromide are due mainly to the fact that it is still less stable 
and more easily oxidized than the latter. It may be made by 
the direct combination of the elements. The action is, however 
slow, reversible, and far from complete. It is catalyzed by plati- 
num, but this simply brings about the equilibrium quicker with- 
out changing the composition of the equilibrium mixture. At 
400C., 80 per cent, of a mixture of equal volumes of hydrogen 
and iodine gases will be converted into hydrogen iodide, or 20 
per cent, of pure hydrogen iodide will be decomposed; and this 
gaseous mixture will remain unchanged in composition so long 
as it is kept at this temperature. If increasing pressure is put 
upon this system, no change will follow because as shown in the 

H 2 + I 2 <=* 2HI 

there is no change in volume as the reaction proceeds. 

As mentioned above, the preparation of hydrogen iodide by 
the action of sulfuric acid upon sodium iodide is excluded be- 
cause of the oxidizing action of the sulfuric acid. By using 
phosphoric acid instead of sulfuric, hydrogen iodide may be pre- 
pared by this method. 

It may be made by the action of iodine upon hydrogen com- 
pounds. An aqueous solution of hydrogen iodide may be easily 
obtained by passing hydrogen sulfide into a suspension of iodine 
in water: 

H 2 S + I 2 = 2HI + S 

The most common method for the preparation of hydrogen 
iodide is by the interaction of iodine, phosphorus and water. 


Phosphorus tri-iodide is first formed and then decomposed by 
water giving hydrogen iodide and phosphorous acid: 

2P + 3I 2 = 2PI 3 
PI 3 + 3H 2 = SHI + H 3 PO 3 

This is analogous to the best method for the preparation of 
hydrogen bromide. 

Physical Properties. Hydrogen iodide fcs a colorless gaseous 
substance much heavier than air as its molecular weight is 128. 
It may be condensed to a slightly yellowish liquid, boiling at 
36C. and freezing at 51C. Hydrogen iodide is very 
soluble in water; the solution saturated at 0C. contains 200 
parts by weight of hydrogen iodide to 100 parts of water. Like 
hydrogen chloride and bromide, it forms a maximum boiling 
mixture with water which boils at 128C. and contains 57 per cent, 
hydriodic acid. 

Chemical Properties. Hydrogen iodide is the least stable of 
the hydro-halides and is so easily oxidized that it is an excellent 
reducing agent. It burns in oxygen, reacts with chlorine with a 
flash of light and the formation of hydrogen chloride and free 
iodine. The liquid is a non-conductor of electricity and is inac- 
tive chemically. But the solution in water, hydriodic acid, is an 
excellent conductor of electricity and is a strong acid, being as 
active as hydrochloric acid solution containing an equal number 
of gram moles per liter. When a colorless solution of hydriodic 
acid is exposed to the air it takes up oxygen, quickly turns 
brown, and finally deposits crystalline iodine. The equation is 
as follows : 

4HI + 2 = 2H 2 + 2I 2 . 

It is this reaction which gives the iodine in the rare cases men- 
tioned above in which it occurs free in nature. 

Oxygen Compounds of Iodine. The oxygen compounds of 
iodine are, in general, much more stable than those of bromine. 
They are formed in analogous ways to the corresponding com- 
pounds of chlorine. The halogen first reacts with water and a 
base to give an iodide and hypoiodite. The hypoiodites are 
even more unstable than the hypobromites, and very quickly 
pass over into the iodides and iodates. 


It will be recalled that iodine occurs in nature largely as sodium 
iodate in the Chile saltpeter beds. 

lodic acid may be prepared from iodates by the second general 
method, but is more readily obtained by the action of nitric acid 
upon iodine, as shown in the following equation: 

I 2 + 10HN0 3 = 2HIO 3 +, 10NO 2 + 4H 2 O 

N0 2 is called nitrogen peroxide. 

lodic acid is a stable substance which crystallizes from its 
very concentrated solutions in lustrous crystals which are ex- 
tremely soluble in water. The solution is a strong acid. By 
carefully heating the crystals, they decompose into water and 
iodine pentoxide, I 2 5 , 

2HI0 3 <= H 2 O + I 2 O 5 

The action is reversible so that the pentoxide is also called iodic 
anhydride. lodic acid is a fairly strong oxidizing agent and 
rapidly reacts with hydriodic acid to form iodine. 

HI0 3 + 5HI = 3H 2 + 3I 2 

Periodates and Periodic Acid. When sodium hypochlorite in 
alkaline solution acts upon sodium iodate, rather difficultly 
soluble sodium periodate, Na 2 H 3 I0 6 , separates out: 

NaI0 3 + NaCIO + NaOH + H 2 = Na 2 H 3 I0 6 + Nad 
This is the disodium salt of orthoperiodic acid, H 5 I0 6 , which may 
be easily prepared by the second general method for the prepa- 
ration of acids. As will be seen, this does not correspond very 
closely with perchloric acid, HC1O 4 . By long and careful 
heating to 100C. in a vacuum the ortho acid, HsIOe, loses water 
and passes over into the meta acid, HKX. 

Oxides of Iodine. Iodine forms two oxides, the tetroxide, 
I 2 4 , and the pentoxide, I 2 5 ; the latter is iodic anhydride. 


The atomic weight of fluorine is 19, which is smaller than 
that of chlorine, and hence, logically, it should have been con- 
sidered first; but, as said before, its properties are not as typical 
of the halogens as are those of chlorine, and consequently it was left 
until this point. 


Occurrence. Fluorine never occurs free in nature, in fact it is 
one of the most difficult of the elements to liberate. It is gener- 
ally found as calcium fluoride (fluor spar), CaF2. Two other 
fairly common compounds are cryolite, NasAlFe, and apatite, 
Ca 5 (PO 4 ) 3 F. These are all difficultly soluble, while the most 
abundant of the naturally occurring compounds of the other 
halogens have been easily soluble. 

Preparation. As we passed from chlorine through bromine 
to iodine, it was noted that it became easier and easier to oxidize 
the compounds, and so to liberate the elements, as the atomic 
weight increased. This would lead us to expect that fluorine, 
the halogen with the smallest atomic weight, would be the 
hardest to oxidize, and this is found to be true; in fact there is no 
oxidizing agent powerful enough to set the fluorine free. There 
remains then only one method for its preparation, electrolysis 
of a compound. Water must be absent because fluorine and 
water immediately react with the formation of hydrogen fluoride 
and ozone, 

3F 2 + 3H 2 O = 3H 2 F 2 + O 3 

Anhydrous liquid hydrogen fluoride is almost a non-conductor 
of electricity, but it dissolves . potassium acid fluoride readily; 
this salt dissociates into ions and the solution becomes a con- 
ductor. Fluorine is given off at the anode and hydrogen at the 
cathode. Possibly potassium is first liberated, and then reacts 
with the hydrogen fluoride, 

H 2 F 2 + 2K = H 2 + 2KF 

The electrolysis is carried out in a copper U-tube cooled to 23 
using platinum-iridium electrodes. 

Physical Properties. Fluorine is a gas having a greenish- 
yellow color much like that of chlorine only less intense. It 
liquefies at low temperatures, boils at 186 and freezes at 223. 
Because of the very great activity of fluorine, it is difficult to 
determine its density; but the molecular weight evidently is not far 
from 38, and since the atomic weight is 19 the formula is F 2 . 

Chemical Properties. Fluorine combines with all the elements 
except oxygen, nitrogen, chlorine and the members of the argon 
group. Hydrogen fluoride is an especially stable compound and 
is instantly formed when fluorine is brought in contact with hydro- 


gen even at a temperature of 252 or with almost any hydrogeri 
compound at ordinary temperatures. The reaction with water 
has been noted above. Fluorine and hydrogen chloride react for 
the liberation of chlorine; of course, it will react even more readily 
on hydrogen bromide and hydrogen iodide. 

Hydrogen Fluoride. Hydrogen fluoride containing a little 
water may be easily prepared by distilling a fluoride, usually that 
of calcium, with sulfuric acid: 

CaF 2 + H 2 S0 4 = CaS0 4 + H 2 F 2 

Here, of course, there is no difficulty caused by the oxidation of 
the hydrogen fluoride. But there is a complication due to the 
fact that hydrogen fluoride will readily attack glass, and hence 
the operation must be carried out in platinum or lead retorts. 
The anhydrous acid may be easily prepared by heating acid 
potassium fluoride in a platinum retort. 

2KHF 2 <=2KF + H 2 F 2 

Physical Properties. Hydrogen fluoride is a clear colorless 
liquid boiling at 19. This is exceptionally high when compared 
with hydrogen chloride, etc. It dissolves readily in water, and 
like the other hydrogen halides forms a constant boiling mixture 
with water. This contains 35 per cent, hydrogen fluoride and 
boils at 120C. The density of hydrogen fluoride vapor, at tem- 
peratures but little above its boiling-point, shows a molecular 
weight of 40, corresponding to H 2 F 2 ; at higher temperatures the 
density decreases, and from the boiling-point of water upward 
the molecular weight is 20 and the formula HF. The existence of 
the compound H 2 F 2 is an important point of difference from the 
other hydrogen halides. 

The formula H 2 F 2 indicates a dibasic acid ; that is, one a mole of 
which can neutralize two moles of sodium hydroxide. This 
is confirmed by the existence of potassium acid fluoride, KHF 2 , 
and others of the same type. The neutral salt, K 2 F 2 or KF, is 
also well known. 

The most interesting property of hydrogen fluoride is its action 
on glass and other silicon compounds. When it acts on glass 
which contains silicates of sodium and calcium, gaseous silicon 
fluoride, SiF 4 , and fluorides of sodium and calcium together with 
water are formed: 



Use is made of this for etching designs on glass. The glass is 
first coated with wax, the design is scratched through, and then 
the whole exposed to the vapor or a solution of hydrogen fluoride. 
Because of this action, it is of course impossible to keep hydro- 
fluoric acid in glass, but it may be kept in wax or gutta-percha 
bottles. The commercial acid is stored and shipped in lead 

The action of hydrofluoric acid on glass is not due to its being 
an especially strong acid, because when we come to examine its 
action upon metals we find that a solution containing 20 grm. 
of hydrofluoric acid in 1,000 grm. of water will act much more 
slowly on a given piece of zinc than will a solution containing 
36.47 grm. of hydrochloric acid in the same amount of water; 
and yet each solution contains the same amount of hydrogen 
which is capable of being replaced by a metal. The electrical 
conductivity of a normal hydrofluoric acid solution is much 
smaller than that of normal hydrochloric acid and this indicates 
that hydrofluoric acid is much weaker than hydrochloric. 

The following table gives in a condensed form a summary of 
the more important of the properties of the halogens. 

Formula of element 

F 2 


Br 2 

I 2 , L 

Atomic weight 





Physical state at ordinary tem- 



+ 59 


+ 184 


light yellow- 


- 7.3 

+ 114 

Oxidizing agent 

ish green 



Hydrogen compounds 

H 2 F 2 -HF 




Physical state. . . . 



+ 19 


68 . 5 



92 3 









Solution in water 
Action toward oxidizing agents. . . 

Heat of formation Gaseous 
Heat of formation dissolved 
Heat of Neutralization NaOH . . . 

weak acid 


strong acid 
oxidized by 

strong acid 
rather easily 

strong acid 
very easily 
+ 55KJ. 






This table shows that there are regular changes in a number of 
the properties of these elements as the atomic weight increases. 
For example the boiling- and melting-points of the elements 
rise as the atomic weight increases. The student will gain 
considerable information if he follows the other properties of 
these elements through in the same way. 

Hydrogen fluoride has an abnormally high boiling-point. In 
this connection, it is significant that it alone of the hydrohalides 
ionizes solutes and that other ionizing solvents boil higher than 
their inactive analogues. 

Hydrolysis. :Salts of weak acids or of weak bases when in 
solution in water act in a very peculiar manner. Their solutions 
are not neutral as are those of the salts of strong acids and bases, 
but the salts of weak acids give alkaline, while those of weak 
bases give acid reacting solutions. Since the salts themselves 
do not contain either the hydroxyl or the hydrogen ions, which 
give the solutions their reactions, the question of the source of 
these ions is one which calls for some little explanation. Evi- 
dently they must have come from the water. But why should 
the hydroxyl ion appear in a solution of a salt of a weak acid, 
and the hydrogen ion in thevsalt of a weak base when neither is 
prominent in a solution of a salt of a strong acid and base? 
The theory of electrolytic dissociation furnishes a satisfactory 

Pure water is a very poor conductor of electricity but it is not 
an absolute non-conductor. Its conductivity has been measured 
and indicates that water itself is dissociated to a slight extent so 
that 1 grni. of hydrogen and 17 grm. of hydroxyl as ion are 
present in eleven million liters of water. Expressing the concen- 
tration in terms of the number of gram ions per liter, the concen- 
tration of each of these substances is the same and they balance 
each other in their action on litmus. 

Reactions between substances in the ionic condition tend to* 
take place in such ways that the concentrations of the ions are 
reduced, either through the formation of substances which pass 
out of solution or those which are but slightly dissociated. 

When a salt of a weak acid and a strong base, say sodium hypo- 
chlorite, is dissolved in water, there are present side by side 
sodium, hypochlorite, hydrogen, and hydroxyl ions. Hypo- 


chlorous acid is a very weak acid, and hence, but slightly dis- 
sociated, while sodium hypochlorite is highly dissociated. There 
is then a chance for a decrease in the ions by the union of the 
hydrogen and the hypochlorite ions, and this they are assumed 
to do to a certain extent. The solution would then contain 
sodium as ion, some hypochlorite, enough hydroxyl to make up 
for the hypochlorite ion which combined with hydrogen as ion, 
and undissociated hypochlorous acid. The undissociated hypo- 
chlorous acid of course does not redden litmus, while the hydroxyl 
ion turns litmus blue. So the solution reacts alkaline and it is 
evident that a salt of any other weak acid should, act in the same 

When the hydrogen ion combines with the anion of the weak 
acid, not only is the hydroxyl left unbalanced in its action upon 
litmus, but the concentration of the hydroxyl increases. This 
comes from more of the water's breaking up in an effort to keep up 
the concentration of the hydrogen. The equation for the 
ionization of water is 

H 2 0<=H+ + OH- 
a b c 

Letting a, b and c represent the concentrations of the water, 
hydrogen and hydroxyl ions respectively, the Mass law equation 


- = constant 

Now a is constant because it stands for the concentration of a 
homogeneous liquid, and hence the product be is constant and 
as b decreases c must increase. 

Salts of weak bases and strong acids react acid because of the 
combination of the cation of the salt with the hydroxyl of the 
water with the formation of the undissociated base, thus leaving 
the hydrogen as ion in excess. 

The phenomena which we have just been discussing are in- 
cluded under the name of hydrolysis. Speaking in terms of the 
ionic theory we may say that hydrolysis consists in the formation 
of an undissociated acid or base by the action of water upon a 
salt of a weak acid or base. 


Solutions of salts of strong acids and bases are neutral, because 
under these conditions there is no tendency for either hydrogen 
or hydroxyl to combine with an ion of the salt; so these salts do 
not hydrolyze. 


The formulas of some of the principal hydrogen compounds of 
the non-metallic elements are given in the following table: 

HF H 2 O H 3 N H 4 C 

HC1 H 2 S H 3 P H 4 Si 

HBr H 2 Se 

HI H 2 Te 

Several of these compounds have not been studied as yet, but 
they are introduced here simply to show the different types of 
combination. H 2 S is hydrogen sulfide, H 2 Se hydrogen selenide, 
H 2 Te hydrogen telluride, HsN or NH 3 ammonia, and HaP 
phosphine, H 4 C or CH 4 marsh gas, H 4 Si silicon hydride. 

From this table it may be seen that these compounds fall 
naturally into four groups according to the number of atomic 
weights of hydrogen which unite with one atomic weight of the 
other element. Each element is considered to have a certain com- 
bining power or valence which is measured by the number of atomic 
weights of hydrogen or of an equivalent substance with which one 
atomic weight of that element can unite to form a compound. This 
number is called the valence of the element. From the definition 
of valence, it follows that in the compounds given in the above 
table, fluorine, chlorine, bromine, and iodine are monovalent; 
oxygen, sulfur, selenium and tellurium divalent; nitrogen and 
phosphorus trivalent; while carbon and silicon are tetravalent. 
Hydrogen is our standard substance for measuring valence and is 
taken to have a valence of one or to be monovalent. Sometimes 
an element does not combine with hydrogen, and then of course 
it is not possible to work out the valence from the hydrogen com- 
pound, and in this case we must resort to some other method. 
One of these is to take a compound with some equivalent sub- 
stance. An equivalent substance is a substance one atomic weight 
of which can unite with, replace, or be replaced by, one atomic weight 
of hydrogen. 


An atomic weight of sodium replaces one atomic weight of 
hydrogen in the compound hydrogen chloride forming sodium 
chloride, NaCl, and is therefore monovalent; also one atomic 
weight of silver is replaced from silver chloride, AgCl, by one 
atomic weight of hydrogen, re-forming hydrogen chloride, there- 
fore, silver is monovalent. Chlorine then in chlorides is deter- 
mined to be monovalent: directly because it was originally com- 
bined with one atomic weight of hydrogen in hydrogen chloride; 
indirectly because it is combined with one atomic weight of the 
monovalent substances silver and sodium. Because chlorine in 
chlorides is monovalent, phosphorus in phosphorus trichloride, 
PCls, is trivalent, while in the pentachloride, PCU, it must be 
pentavalent. It is not necessary, however, that we take a com- 
pound with a substance exactly equivalent to one atomic weight 
of hydrogen; for once we have established the valence of any 
given element, it may in turn be used to measure that of others. 
One atomic weight of oxygen is united with two atomic weights 
of hydrogen in the compound water, hence oxygen is divalent 
and is commonly used as a standard for determining valence. 
Take for example, phosphorous trioxide, P^e; the six atomic 
weights of oxygen have together a valence of twelve, which is 
divided between the four atomic weights of phosphorus, showing 
that each atomic weight of the latter has a valence of three. In 
the pentoxide, P2O 5 , the valence of the five oxygens is ten, which 
is divided between the two of phosphorus; and therefore in this 
compound the phosphorus has a valence of five, as in the 

The examples just given show that the valence of an element 
is not a fixed constant. However, the number of valences 
belonging to any one element is generally small. Nitrogen has 
an unusually large number of valences, as may be seen from the 
following list of its oxides, N 2 O, NO, N 2 3 , N0 2 , N 2 5 . In these 
it varies from monovalent to pentavalent. Comparatively few 
elements have more than two valences so that these properties 
are not difficult to learn, and they are of great assistance in the 
formulation of compounds. 

Equivalent Weights. The conception of valence enables us 
to get easily another useful conception, that of equivalent 
weight. This is the number of parts by weight of an element 


which will combine with, replace or be replaced by 1.008 parts 
by weight of hydrogen, or 8 parts of oxygen. Or it may be 
defined as the atomic weight of the element divided by its 
valence. For example, the equivalent weight of sulfur in 
hydrogen sulfide is 32 -^; of nitrogen in ammonia is 14>0 J^; 
of carbon in marsh gas is *%. From the second form of the 
definition of equivalent weight, it is evident that an element has 
as many equivalent weights as it has valences. 

These equivalent weights are like the atomic weights in many 
respects, except that they represent directly without any multiple 
the proportions by weight in which the elements combine. 



General. Sulfur is an element which has been known since 
prehistoric times. It occurs fairly abundantly in nature, both 
free and in compounds. The free sulfur is generally found in 
volcanic regions, but several deposits of non-volcanic sulfur are 
known, notably those of Louisiana which now furnish most of 
the sulfur used in the U.nited States together with large quan- 
tities for export. In combination, sulfur occurs chiefly in the 
form of sulfates, such as gypsum, CaSO4-2H 2 O j Glauber salts, 
Na2SO 4 -10H 2 O, and Epsom salts, MglSO^THg^ and as sulfides, 
zinc blende, ZnS, and galena, PbS. It is present in 

combination in many portions of the animal and plant organisms. 

In many of its chemical properties, sulfur shows a rather close 
relationship to oxygen. Like oxygen it is able to assume allo- 
tropic forms, and is somewhat remarkable for the large number of 
such modifications which are known to exist, two and prob- 
ably more crystalline modifications, two liquid, and at least two 
gaseous forms. 

Extraction of Sulfur. Most of the sulfur from volcanic sources 
comes from Sicily where it is found in large quantities mixed 
with limestone, gypsum, celestite and other minerals. The sulfur 
is extracted from the ore by heating in oven-like kilns or cells 
made of masonry. The greater part of the heat is furnished by 
burning coke which is mixed with the ore and charged into the 
cells. Of course, part of the sulfur burns. Economy of fuel is 
secured by arranging the cells in batteries of two or more and 
passing the gases from the one in action into a freshly charged 
cell in order that the waste heat from the first may warm the 
charge in the second. The sulfur melts and runs down to the 
floor of the cell from which it is drawn off and cast into conical 
wooden moulds forming the roll sulfur or brimstone of commerce. 
In some more recent installations the sulfur is extracted by melt- 




ing it out with high-pressure steam. This method is being used 
with success in some rather recently developed mines in Mexico. 

In Louisiana, the sulfur is obtained by drilling down about 
500 ft. to the sulfur-bearing deposits and forcing down into the 
holes water heated to a temperature of 168 which is far above 
the melting-point of sulfur (120). The water must, of course, be 
under pressure of about seven atmospheres to reach this tempera- 
ture. The sulfur melts and collects at the bottom of the hole; 
from there it is pumped out by compressed air. The sulfur can 
be obtained in this way very cheaply, and is more than 99.5 
per cent. pure. A single well often yields 500 tons of sulfur per 
day and the entire mine has produced 122,000 tons in two months 
which is more than the world's consumption for that period. 
The production of sulfur in the United States in 1914 was 328,000 
long tons. 

Purification of Sulfur. When it is necessary to purify roll 
sulfur, the process of distillation is employed. The sulfur is 
boiled in iron retorts and the vapor is condensed in large brick 
chambers. At first when the walls of the chamber are cool, the 
sulfur condenses in the form of a rather fine powder, and is put 
on the market under the name of flowers of sulfur. When the 
chamber becomes heated from the 
heat of condensation of the sulfur, 
the latter collects in the form of a 
liquid and is drawn off into moulds 
as before. 

Crystalline Forms of Sulfur. Crys- 
tals of rhombic (Fig. 30) or a sulfur 
are found in nature and may be easily 
prepared by dissolving sulfur in car- 
bon disulfide, in which it is readily 
soluble, and allowing the solution to evaporate slowly. The 
sulfur is deposited in octahedral-like crystals which have a 
density of 2.07, and melt to a clear yellow liquid at 113. If 
the clear crystals are kept for some time at temperatures 
between 100 and the melting-point they gradually become 
opaque and brittle, and when this change is complete, will melt 
at 120 instead of 113, and have a density of 1.96 instead of 
2.07. They then have undergone a transformation into or 




FIG. 30. 

I T 



monoclinic sulfur (Fig. 30). This change is hastened by 
scratching the rhombic crystals with some hard point or by 
simply touching them with a monoclinic crystal. 

Well denned crystals of monoclinic sulfur are easily obtained 
by melting some sulfur and allowing it to cool until a crust has 
formed on top. A hole is then made through this, and the still 
molten sulfur is poured out. The sides and bottom of the vessel 
will then be found covered with long needle-shaped monoclinic 
crystals. When first formed the crystals are amber colored, and 
may be bent slightly without breaking. If the crystals are kept 
at 100 or above, they retain their properties indefinitely, but 
if kept at ordinary temperatures for a day or so they change 
color to the lemon yellow of ordinary sulfur, become opaque, and 
increase in density to 2.07, in short they change into a or rhom- 
bic sulfur. This transformation is hastened by contact with 
rhombic crystals. The external form of the crystals is that 
of the monoclinic, while internally they consist of an aggregation 
of rhombic crystals. Such crystals as these, which are internally 
of one form and externally of another, are called pseudomorphs. 

Transition Point. The facts just given indicate that the mono- 
clinic form is stable at higher temperatures while the rhombic is 
at lower. There must be then a temperature at which they are 
both stable. This temperature is 96 and is called the transition 
point of these modifications of sulfur. From 96 to 120 the 
monoclinic modification is stable while the rhombic is stable 
below 96. 

At 96 the two modifications are in equilibrium, and hence 
have the same vapor pressure, solubility, etc. The monoclinic 
can exist below 96, but in an unstable state and under these con- 
ditions it is more soluble than the rhombic and has a higher vapor 
pressure. Correspondingly, above 96 the rhombic is the un- 
stable form and has the greater solubility and vapor pressure. 
These are simply examples of the general rule that the less stable 
modification has the higher solubility and vapor pressure. 

Although unstable, the monoclinic sulfur can exist for a time 
below 96 and the rhombic above this temperature. This is 
due to what is called suspended transformation and is much 
the same phenomenon as the supercooling or superheating of 


Liquid Sulfur. Monoclinic sulfur melts at 120 to a light 
yellow, mobile liquid, which on cooling passes back into mono- 
clinic crystals. If, however, the temperature be raised to 160 
the liquid suddenly becomes dark in color and so viscous that the 
vessel may be inverted without the contents running out. At 
180 the maximum viscosity is reached; at 260 the viscosity 
begins to decrease noticeably, and at 444.6, the boiling-point of 
sulfur, the liquid is mobile once more. 

The cause of this remarkable change at 160 is the appearance 
of a new form of liquid sulfur called /z sulfur. When the dark 
colored liquid is slowly cooled the reverse change takes place, and 
the light yellow form reappears. 

Amorphous Sulfur. By heating sulfur to temperatures near 
its boiling-point and suddenly cooling the mass by pouring it 
in a thin stream into water, a clear, transparent, plastic, semi- 
fluid substance is obtained, which is called plastic sulfur. After 
a few days it becomes hard and opaque and is then found to con- 
sist of a mixture of rhombic crystals, soluble in carbon disulfide, 
and another form of sulfur which is insoluble in carbon disulfide. 
It is not crystalline and hence is called amorphous sulfur. It is 
supercooled sulfur, and changes slowly into the crystalline variety . 
The change takes years for completion at ordinary temperatures, 
but is finished in an hour at 100. 

There is another amorphous modification, soluble in carbon 
disulfide, and also soluble in water. 

Molecular Weight. At 444.6 sulfur boils. The vapor has an 
orange yellow color when just above its boiling-point. At 500 
it is a deep red and then becomes lighter until at 650 it is a light 
yellow. The molecular weight at temperatures near the boiling- 
point is 230 and gradually falls until at 1,000 it reaches 64 and 
remains constant at this figure as the temperature is raised. 
Since the atomic weight of sulfur is a trifle over 32, the molecular 
weight of 230 corresponds to a formula of something be- 
tween S 7 and S 8 , while that of 64 indicates a formula of S 2 . 
The molecular weight of the element in solution as determined 
from the lowering of the freezing-point shows that when dis- 
solved it has the formula S 8 . 

Chemical Properties of Sulfur. Most all metals when finely 
powdered and rubbed in a mortar with sulfur combine directly 


with the latter to form sulfides. At higher temperatures, the 
combination takes place so rapidly and with the evolution of so 
much heat that the general phenomena of combustion, the pro- 
duction of light, etc., are exhibited. The element will also 
unite with many of the non-metals; the halogens, hydrogen, 
and oxygen for example. 

With oxygen it forms the two very important compounds, 
sulfur dioxide, S0 2 , and sulfur trioxide, S0 3 ; these are the anjjj:- 
dricles of sulfurpus and sulfuric acids respectively. With hy- 
drogen it forms the important compound hydrogen sulfide, H 2 S. 

Uses of Sulfur. Thousands of tons of sulfur are used in the 
manufacture of sulfur dioxide which in turn is used for making 
sulfuric acid, in bleaching straw, wool, feathers, etc., and as a 
germicide. Sulfur is also employed in making carbon disulfide, 
gunpowder, matches, fireworks, sprays for fruit trees, and 
vulcanized rubber. 

Hydrogen Sulfide. Hydrogen sulfide is a gaseous substance 
whose formula is H 2 S. It occurs in nature in the gases of some 
volcanoes, and is found in many mineral waters, which are com- 
monly called sulfur waters. It may be prepared by the direct 
union of the elements by passing a stream of hydrogen through boil- 
ing sulfur. The reaction is reversible and hence is not complete. 

Hydrogen sulfide is usually prepared by acting upon a sulfide, 
generally ferrous sulfide, FeS, with either dilute sulfuric or 
hydrochloric acid. The equation is 

FeS + 2HC1 = H 2 S + FeCl 2 

The ferrous sulfide is not very soluble in water but what does 
dissolve is largely broken up into ferrous and sulfur ions. Upon 
the addition of the acid, the sulfur ion unites with the hydrogen 
to form the very slightly dissociated compound hydrogen sulfide, 
H 2 S. The water soon becomes saturated with this and then the 
substance is evolved in the gaseous state. The conditions are 
especially favorable for reaction in this case because we have a 
substance formed which is both slightly dissociated, and only 
moderately soluble. As fast as the sulfur ions present in solution 
are used up, more ferrous sulfide dissolves to keep up the supply. 
So the process continues until either the ferrous sulfide or the 
acid is consumed. 


Hydrogen sulfide is also formed during the decay of albuminous 
substances such as eggs, etc. The odor of such substances in 
this condition is partly due to its presence. 

Physical Properties. Hydrogen sulfide is at ordinary tem- 
peratures a colorless, gaseous substance whose molecular weight 
34.09. Its critical temperature is 100, so it can be liquefied 
by pressure alone at ordinary temperatures. The critical pres- 
sure is 90 atmospheres. It boils at 61 and freezes at 83. 
The liquid is chemically inactive. Hydrogen sulfide has a very 
disagreeable odor much like that of rotten eggs, and acts as a 
powerful narcotic poison when inhaled. Very dilute chlorine 
gas is used as an antidote. 

Hydrogen sulfide is somewhat soluble in water. One volume 
of water dissolves 4.37 volumes at and 3.23 at 15. The 
solubility follows Henry's law that the volume of a given gas dis- 
solved by a fixed mass of a solvent is independent of the pressure, 
while the mass of the gas dissolved is directly proportional to the 

Chemical Properties. Hydrogen sulfide burns in air with a 
pale blue flame, forming water and sulfur dioxide, S0 2 , 

2H 2 S + 30 2 = 2H 2 + 2SO 2 

The equation shows that two volumes of H 2 S require three 
volumes of oxygen for complete combustion. When mixed in 
these proportions the two gases explode violently upon the 
passage of an electric spark. 

When the hydrogen sulfide is burned with an insufficient supply 
of air, water and free sulfur are the main products. This fact is 
taken advantage of for the recovery of sulfur from certain classes 
of waste products which will easily yield hydrogen sulfide. 

With sulfur dioxide, hydrogen sulfide rapidly reacts, forming 
sulfur and water, 

2H 2 S + SO 2 = 3S + 2H 2 O 

This is probably one of the reactions which gives rise to the 
volcanic sulfur. 

When heated to 300 it is slightly decomposed into hydrogen 
and sulfur, and at a white heat the decomposition is practically 
complete. At each temperature, there is a certain mixture of 


hydrogen, sulfur, and hydrogen sulfide, at which equilibrium 

Solution in Water. A solution of hydrogen sulfide in water 
acts slightly acid toward litmus, and hence is sometimes called 
hydrosulfuric acid. Hydrogen sulfide is a very weak acid, and 
is but slightly dissociated. It differs from the acids which we 
have discussed in detail, except hydrofluoric, in that the two 
atomic weights of hydrogen per mole of hydrogen sulfide are both 
replaceable by metals and hence are ionizable. This state of 
affairs is described by the term dibasic, which means that one 
mole of the acids contains two atomic weights of replaceable 
hydrogen, and will react with two moles of a base like sodium 

H 2 S + 2NaOH = Na 2 S + 2H 2 O 

Dissociation of Dibasic Acids. Hydrogen sulfide like all other 
dibasic acids, shows a marked tendency to ionize in such a way 
that only one hydrogen per mole is split off, as shown by the 
following equation, 

H 2 S <= H+ + HS- 

This ionization is especially prominent in concentrated solutions. 
As the solution is diluted, or still better as the concentration of the 
hydrogen ion is reduced by neutralization with a base, the hydros 
sulfide ion r HS~. breaks up into hydrogen and sulfur as ions, 

As a result of this tendency of dibasic acids, it is possible to 
obtain two series of salts of these acids according as the one or 
both hydrogen ions are replaced by metals. For example, we 
have the compounds NaHS and Na 2 S. The compound NaHS is 
often called sodium hydrosulfide from its resemblance to sodium 
hydroxide, NaOH, while Na 2 S is called sodium sulfide. Several 
other names for the same compounds are in common use as 
shown by the following table: 


Acid sodium sulfide Neutral sodium sulfide 

Primary sodium sulfide ' Secondary sodium sulfide 

Mono-sodium sulfide Di-sodium sulfide 


The terms acid and neutral sodium sulfide are rather deceptive, 
since the so-called acid salt is really -just about neutral in its 
action upon litmus, while the neutral salt is decidedly alkaline. 
This latter effect is due to hydrolysis which changes the neutral 
sulfide nearly completely to the acid sulfide according to the 

Na 2 S + H 2 O <= NaOH + NaHS 

2Na+ + S- - + H+ + OH- <= 2Na+ 

Analytical Reactions of Hydrogen Sulfide. The sulfides of 
many of the metals are very slightly soluble in water and are 
consequently precipitated when hydrogen sulfide is passed into 
aqueous solutions of their salts. For example, a solution of 
mercuric chloride, HgCl 2 , gives a black precipitate of mercuric 
sulfide, HgS, with hydrogen sulfide, 

HgCl 2 + H 2 S = HgS + 2HC1 

Hg+ + + 2C1- + 2H+ + S- ~ < HgS + 2H+ + 2C1- 

Salts of copper, lead, bismuth, antimony, tin, and several other 
metals act in the same way, and their sulfides are so difficultly 
soluble that they will be formed even in fairly acid solutions. 

Other sulfides are more soluble, and require a higher concen- 
tration of sulfur as ion for their formation. This can be obtained 
by taking advantage of the fact that salts of weak acids are 
highly dissociated, and using instead of hydrogen sulfide, sodium 
or ammonium sulfide. For example, ferrous sulfide is not 
precipitated by hydrogen sulfide, but is by sodium sulfide, 

Fe+ + + 2C1- + 2Na+ + S~ - <=* FeS + 2Na+ 4- 2C1~ 

Use is made of these properties of the sulfides to separate the 
metals into groups according to similarities in their behavior, 
and this together with the fact that many of the sulfides have 
distinctive colors, makes hydrogen sulfide an exceedingly im- 
portant reagent. 

Hydrogen Sulfide as a Reducing Agent. An aqueous solution 
of hydrogen sulfide when exposed to the air loses its odor and 
becomes milky, finally depositing a white precipitate of finely 


divided sulfur. This is formed by the oxidation of the hydrogen 
sulfide by the oxygen of the sir, 

2H 2 S+ 2 = 2H 2 + 2S 

The oxygen oxidizes the hydrogen sulfide, and the latter reduces 
the oxygen. 

When hydrogen sulfide is passed into water in which iodine is 
present a solution of hydriodic acid is obtained together with 
free sulfur, 

I 2 + H 2 S = 2HI + S 

or expressed as ions, 

I 2 + 2H+ + S- - = 2H+ + 21- + S 

In this case the action is easily seen to be between the sulfur 
ion and the thing reduced the iodine. Upon examination of 
the other cases in which hydrogen sulfide acts as a reducing agent, 
it will be found that in the majority of cases the action consists 
in the transformation of the sulfur ion into free sulfur, but some- 
times sulfuric acid is formed. The hydrogen has nothing to do 
with the reducing action. 

Hydrogen sulfide will reduce concentrated sulfuric acid to 
water and sulfur dioxide, sulfur being set free, 

H 2 S + H 2 S0 4 = S + 2H 2 + S0 2 

Because of this action, concentrated sulfuric acid cannot be used 
to dry hydrogen sulfide. 

Polysulfides. When a solution of a soluble sulfide or acid 
sulfide is brought in contact with sulfur, the latter is dissolved 
and the solution becomes dark yellow in color. If sodium sul- 
fide is used, compounds Na 2 S 4 , and Na 2 Ss may be obtained. 
These substances are called polysulfides. They are salts and 
yield ions which are Na + and anions which vary in composition 
from S 4 or HS 4 ~ to S 5 or HS 5 ~. 

Sulfur Dioxide or Sulfurous Anhydride. Sulfur dioxide, at 
ordinary temperatures, is a gaseous substance which is present 
in the gases given off by many volcanoes. It is formed, as we 
have seen, by burning sulfur or hydrogen sulfide in the air. It 
may also be prepared by roasting metallic sulfides, pyrite, FeS 2 , 


for example, in the air. When these are raised to the required 
temperature they burn, forming sulfur dioxide, S0 2 , and the 
oxides of the metals, 

4FeS 2 + 11O 2 = 2Fe 2 3 + 8S0 2 

Many of the metals are obtained from their sulfides and such 
a roasting process is generally a necessary preliminary to the 
reduction to the metal. 

Sulfur dioxide may be very conveniently prepared by the 
action of acids upon sulfites or acid sulfites. Under these 
conditions we should expect the formation of sulfurous acid, 
H 2 SO 3 , and this is doubtless formed first but breaks down into 
water and SO 2 . The most convenient laboratory method for 
its preparation is based upon this, sodium acid sulfite, NftTISfjJU 
and strong sulfiiric acid being used: 

2NaHSO 3 + H 2 SO 4 = Na 2 S0 4 + 2S0 2 + 2H 2 O 

It is often prepared by the action of reducing agents upon 
concentrated sulfuric acid. The most suitable ones are perhaps 
metallic copper or free sulfur, the equations being: 

2H 2 SO 4 + Cu = CuSO 4 + SO 2 + 2H 2 O 

2H 2 SO 4 + S = 2H 2 O + 3SO 2 

A good deal of sulfur dioxide is prepared for technical purposes 
by acting upon concentrated sulfuric acid with carbon, 

2H 2 SO 4 + C = 2H 2 O + CO 2 + 2SO 2 

The presence of the carbon dioxide mixed with the sulfur dioxide 
usually makes this method unsuitable for laboratory purposes. 
The greater part of the sulfur dioxide which is used in the arts 
is made by burning sulfur or pyrite. 

Physical Properties. Sulfur dioxide is a colorless gaseous 
substance with a strong odor which is familiar as the odor of 
burning sulfur. Its critical temperature is 156 and the critical 
pressure is 79 atmospheres. It may, then, be. liquefied by pressure 
alone at ordinary temperatures. Since its boiling-point is 10, 
it may be easily liquefied by a mixture of ice and salt. The liquid 



is as clear and colorless as water. It is a fairly good solvent and 
is able to ionize many salts, though to a smaller extent than water. 
Liquid sulfur dioxide is now an article of commerce, being sold 
in metal cylinders from which it may be removed as desired. 

At ordinary temperatures sulfur dioxide deviates widely from 
the gas laws a fact which is connected with its ease of lique- 
faction. It is rather soluble in water, one volume of the latter 
taking up at ordinary temperatures 50 volumes of sulfur dioxide 
and a smaller quantity at higher temperatures. At higher 
temperatures the solubility approximately follows Henry's 
law. The aqueous solution smells strongly of the dioxide which 
can be completely removed by boiling since it does not form 
either a maximum or minimum boiling mixture with water. 

Chemical Properties. Sulfur dioxide is stable and is decom- 
posed only at very high temperatures. It will unite directly 
with chlorine for the formation of sulfuryl chloride", S0 2 C1 2 , or 
^with oxygen to form sulfur trioxide, S0 3 . The solution of 
sulfur dioxide in water reacts acid and when neutralized with 
^sodium hydroxide and the solution evaporated, it yields the 
salt, sodium sulfite, Na 2 SO 3 . We therefore say that the solu- 
tion contains sulfurous acid, and ascribe to this the formula 
JI 2 SO 3 . A pure acid of this composition cannot be prepared 
since it breaks down into water and sulfur dioxide as the solu- 
tion is concentrated. Because of the formation of sulfurous 
acid, sulfur dioxide is often called sulfurous anhydride. 

Like hydrogen sulfide, sulfurous acid is a dibasic acid and 
yields acid and normal salts. It is a stronger acid than hydro- 
gen sulfide, and the normal salts are not so highly hydrolyzed 
as the normal sulfides. The hydrolysis, however, is marked. 

Reducing Action. Sulfurous acid and the sulfites are strong 
reducing agents, passing into sulfuric acid or sulfates. As an 
example, we may take the reaction of sulfurous acid upon chlorine 
and water, 

H 2 S0 3 + C1 2 + H 2 = H 2 SO 4 + 2HC1 

Here the oxygen necessary for the formation of sulfuric acid 
comes from the water. The sulfurous acid may be replaced 
by sulfites. 

Because of this reaction between chlorine and sulfurous acid, 


sulfurous acid and sulfites are used as "antichlors" in the bleach- 
ing industries to remove the last of the chlorine. 

Sulfites are spontaneously oxidized by the oxygen of the 
air, an action which may be greatly slowed down by the addi- 
tion of small quantities of sugar, alcohol, or glycerine, which 
act as negative catalyzers. 

Bleaching Action. Sulfurous acid and its salts are rather 
powerful bleaching agents and are extensively used for bleach- 
ing silk, wool, feathers, straw, etc., which would be destroyed 
by chlorine or hypochlorous acid. The sulfurous acid acts very 
differently from the chlorine in that it reduces the coloring 
matter instead of oxidizing it. The coloring matters are not 
destroyed in this case and after the process is finished the goods 
must be carefully washed to remove the altered coloring matter, 
as otherwise the oxygen of the air would soon restore the color. 
If sulfur dioxide is used, the goods are first moistened so that 
sulfurous acid may be formed. 

Action on Organisms. Sulfur dioxide, sulfurous acid and sul- 
fites are powerful poisons toward vegetable organisms of all 
kinds. Because of this property, sulfur dioxide is used as a 
germicide and fungicide. Sulfurous acid and sulfites have been 
much used as preservatives in food products, but are now gen- 
erally forbidden. 

Sulfur Trioxide or Sulfuric Anhydride. Sulfur trioxide is 
formed by the direct union of sulfur dioxide and oxygen, 

2SO 2 + O 2 ^ 2SO 3 

The action, however, takes place very slowly even at some- 
what elevated temperatures. In this case, one cannot resort 
to raising the temperature to a very high point in order to in- 
crease the rate of combination, because of the fact that at a high 
temperature the action is reversed. However, there are several 
substances, notably finely divided platinum and ferric oxide, 
Fe 2 O 3 , which will actively catalyze the union of the gases. In the 
case of the finely divided platinum, the process takes place to 
the best advantage when the temperature is about 400. The 
sulfur trioxide so formed may be condensed to a colorless mobile 
liquid which boils at 46. When cooled, it forms a transparent 
glassy solid which melts at 15. After standing for some time, 



it passes over into a more stable form which appears as a white, 
opaque, glistening mass of needle-shaped crystals. This opaque 
form does not melt, but at 50 passes directly into the vapor 
which upon being cooled condenses to the liquid, from which 
the ice-like solid may first be obtained and then finally the 
opaque form. The white, crystalline solid has the formula 
S 2 O 6 , while that of the other modification is SO 3 . The S 2 O 6 
is called a polymer of the SO 3 . The formation of the liquid and 
glassy solid before the appearance of the stable opaque form is 
another instance of the phenomena described by the law of 
successive reactions. 

Action on Water. When sulfur trioxide comes in contact 
with moist air, it forms dense white fumes, consisting of sul- 
furic acid formed by the union of the sulfur trioxide with water 
vapor. When the trioxide is brought in contact with water, 
combination takes place with the production of so much heat 
that it hisses as though a red hot body were plunged into the 
water. The result of this action is sulfuric acid, so sulfur 
trioxide is often called sulfuric anhydride. 

Sulfuric acid is an exceedingly important chemical as may be 
seen from its yearly production which amounts to over 2,000,000 
tons in this country alone. It is made by two processes. The 
first, called the contact process, is very simple in theory, but has 
certain practical difficulties, and is best adapted to the pro- 
duction of very concentrated acid. 

The second is the lead chamber process, the chemistry of 
which is more complex but whose operation is sure and simple. 
This turns out immense quantities at a low cost of a somewhat 
dilute acid suitable for many purposes, the bulk of it being 
used for preparing phosphate fertilizer. 

The starting-point in each process is the preparation of sulfur 
dioxide either by burning sulfur in the air or more commonly 
by roasting a sulfide such as pyrite, FeS 2 ; in the latter case the 
reaction is, 

4FeS 2 + 110 2 = 2Fe 2 O 3 + 8SO 2 (1) 

In each process with the aid of catalyzers this sulfur dioxide, 
oxygen of the air and water are then caused to react for the 
formation of sulfuric acid. The chemistry of the contact process 



is very simple. Three reactions in addition to that in equation 
(1) are involved: First the formation of sulfur trioxide, 

2S0 2 + 2 + 2SO 3 (2) 

Second, the union of the sulfur trioxide with sulfuric acid to 
form fuming or pyrosulfuric acid, H 2 S 2 C>7 

S0 3 + H 2 SO 4 * H 2 S 2 0; 


Third, the dilution of this fuming acid with just enough water 
to make the product 100 per cent, sulfuric acid 

H 2 + H 2 S 2 7 = 2H 2 SO 4 (4) 

The reaction between sulfur dioxide and oxygen (2) is the 
troublesome one. It is slow, reversible, and far from complete 
at temperatures above 400. Because of its slowness a positive 
catalyzer, either platinum or ferric oxide, must be used. Plati- 
num works rapidly at 400, and at this temperature 98-99 per 
cent, of the SO 2 may be oxidized to trioxide. 

But the platinum is quickly " poisoned " by the small quantities 
of arsenic compounds present in the dust from the sulfur or 

A = Pyrites Kiln = Drying Tower 

B - Dust Chamber F=- Testing Box 

C = Lead Cooling Pipes G = CpntactChamber 

D Washing Towers 


H= Absorption Vessel 

FIG. 31. 

pyrite burners, and soon loses its activity as a catalyzer unless 
the sulfur dioxide is purified with painful care by first allowing 
the dust to settle (Fig. 31), then cooling the gas, washing it with 
a spray of water and drying with sulfuric acid before it goes to 
the catalyzer in the contact chamber. Ferric oxide is not 
poisoned by arsenic but works rapidly enough only at tem- 
peratures of 625 or higher, and at 625 not more than 70 per 
cent, of the dioxide can be converted into trioxide. Hence 
ferric oxide is much less efficient than the more expensive 


and troublesome platinum, so platinum is the catalyzer most 
commonly used.. The plant is shown diagrammatically in Fig. 
31. The reaction represented by equation (1) takes place 
in A; by equation (2), in G; and (3), in H; where the sulfur 
trioxide, removed from the nitrogen and oxygen left from the air 
by absorption in concentrated sulfuric acid, forms fuming sulfuric 
acid, which upon dilution gives sulfuric acid as shown in (4). 

Lead Chamber Process. In this very important process, 
nitrogen peroxide, N0 2 , catalyzes the formation of sulfuric acid 
from the sulfur dioxide, oxygen and water. 

The sulfur dioxide is prepared as before, see equation (1). The 
oxygen comes from the air. The nitrogen peroxide, NO 2 , comes 
from the action of sulfur dioxide and water upon nitric acid, giv- 
ing sulfuric acid and nitric oxide, NO; the latter then reacts with 
the oxygen of the air to form nitrogen peroxide. The equations 
are as follows: 

3SO 2 + 2HNO 3 + 2H 2 = 3H 2 S0 4 + 2NO (5) 

2NO + 2 <= 2N0 2 (6) 

The water is introduced either in the form of steam or as a spray 
of very finely divided water. The reactions which produce the 
greater part of the sulfuric acid consist in the interaction of the 
sulfur dioxide, oxygen, nitrogen peroxide and water for the for- 

mation of a compound called nitrosylsulfuric acid, S0 2 <f ^~ 

and its subsequent decomposition by water, forming sulfuric 
acid and a mixture of nitric oxide and nitrogen peroxide. The 
equations are 

4S0 2 + 2H 2 + 4N0 2 + O 2 = 4SO 2 < (7) 


2S 2< \N0 2 + H2 = 2H2S 4 + N0 + N 2 (8) 

The nitric oxide then reacts with some more oxygen as shown 
in equation (6) to reform the peroxide which will at once react 
with more of the sulfur dioxide, oxygen, and water to form more 
nitrosylsulfuric acid, "which is decomposed by water, forming 
another lot of sulfuric acid and reproducing the oxides of nitrogen. 



This cycle of changes goes on until all but one-half a per cent, of 
the sulfur dioxide is used up. So a very small quantity of the 
oxides of nitrogen will transform a large amount of sulfur dioxide. 
These reactions take place in very large lead lined rooms or 
chambers (Fig. 32), of which there are three or more in each 
plant. These chambers are often 100 ft. or more in length 
and hold from 150,000 to 200,000 cu. ft. Lead is used in their 
construction because it is but little attacked by dilute sulfuric 
acid. The sulfuric acid formed collects on the bottom of the 

50 t 



FIG. 32. 

chambers, while a mixture of the oxides of nitrogen and the 
nitrogen from the air which furnished the oxygen issues from 
the last chamber. The oxides of nitrogen are too valuable to 
be lost so advantage is taken of the fact that they may be ab- 
sorbed in concentrated sulfuric acid forming nitrosylsulfuric acid, 

NO + N0 2 + 2H 2 S0 4 = 

H 2 


This, it will be noticed, is equation (8) written reversed. The 
absorption takes place in what is called the Gay Lussac tower. 
This is about 50 ft. high and is filled with tiles over which con- 
centrated sulfuric acid runs in a thin stream, while the gases 
from the last lead chamber enter at the bottom and leave near 
the top. The acid, which collects at the bottom of the tower, 
contains the nitrosylsulfuric acid in solution. This may be 



easily decomposed by water, regenerating the oxides of nitrogen 
necessary for catalyzing the formation of the sulfuric acid. 
But to be of service, these oxides of nitrogen must be reintroduced 
into the first lead chamber. This is done by elevating the acid 
drawn off from the bottom of the Gay Lussac tower, with com- 
pressed air to the top of a similar structure called the Glover's 
tower which just precedes the first lead chamber. Here the 
concentrated sulfuric acid is mixed with dilute acid from the 
chambers. This partly decomposes the nitrosylsulfuric acid. 
(See equation 8). The mixture is then allowed to flow slowly 
down through 'the tower over the acid-resisting stones or tiles. 
On its way down it meets the hot sulfur dioxide and air from 
the burners. These gases enter at the bottom of the tower, 
pass up through and out near the top, going from there into the 
first lead chamber. The hot gases while flowing through the 
tower remove the oxides of nitrogen from the diluted sulfuric 
acid and evaporate a large part of the water. This returns 
the oxides of nitrogen to the point where they are needed, con- 
centrates the sulfuric acid, and at the same time cools the gases 
to the temperature favorable for the reactions. 

Other reactions than the ones given above take place during 
the process, whereby a part of the nitrogen compounds are re- 
duced to nitrous oxide, N 2 0, and, as this takes no further part 
in the reaction, it constitutes a loss of active nitrogen to the 
system. To make up for this and for mechanical losses, it is 
necessary to use about 4 parts of sodium nitrate for every 100 
parts of sulfur burned, to make nitric acid enough to replace the 
oxides of nitrogen lost. The nitric acid is usually added at 
the top of the Glover's tower and here reacts with the sulfur 
dioxide as shown in equation (5). 

The whole process may be briefly summed up as follows: 
Sulfur dioxide is formed in the pyrite burner, equation (1). 
The hot gas passes either directly to the Glover's tower or some- 
times first through a dust-settling chamber. In the Glover's 
tower the oxides of nitrogen are removed from the nitrosyl- 
sulfuric acid, equation (8), losses replaced by the addition 
of nitric acid, equation (5), and part of the chamber acid con- 
centrated. In addition, reactions represented by equations (6), 
(7), and (8) take place to a certain extent in the tower. From 

SULFUR . 179 

the Glover's tower, the gases pass on through the lead chambers in 
which the principal reactions are represented by equations (6), (7), 
and (8). From the last chamber, the residual gases pass through 
the Gay Lussac tower where the oxides of nitrogen dissolve in 
the concentrated sulfuric acid, forming nitrosylsulfuric acid 
as shown in equation (9). The nitrosylsulfuric acid then goes 
back to the Glover's tower and the oxides of nitrogen begin 
the journey once more. 

Concentration of the Acid. The chamber acid from the 
sulfuric acid plant runs about 65 per cent, acid and 35 per cent, 
water. Enormous quantities of it are used at this strength, for 
example, in making fertilizer, but for many purposes the greater 
part of the water must be removed. This is very commonly done 
by running the chamber acid together with that from the Gay 
Lussac tower through the Glover towers until the acid has reached 
a concentration of 78 per cent, or higher. It may then be con- 
centrated still further by heating in cast-iron retorts, because 
sulfuric acid of this strength or stronger does not attack iron, 
while the weaker acid does. An entirely adequate explanation 
of this behavior is lacking. 

Stills of glass or of platinum lined with gold are sometimes 
used instead of the cast-iron ones. 

Another method of concentration which is sometimes used 
is to evaporate the chamber acid in lead pans heated from 
above until the acid has a density of 1.71 corresponding to 
78 per cent, acid, when it becomes a strong enough oxidizing 
agent to attack the lead rapidly. Further concentration 
must then be done in cast-iron, glass, or platinum. 

Physical Properties. Commercial sulfuric acid or oil of vitriol 
has a density of 1.83-1.84 and contains 94 per cent, of the acid. 
The strongest sulfuric acid which can be secured by evaporation 
has a density of 1.841 and contains 98.2 per cent, of the acid. 
This is the composition of the maximum boiling mixture. It 
boils at 330 C. The vapor consists largely of a mixture of sulfur 
trioxide and water as may be demonstrated by the fact that 
a flask filled with sulfuric acid vapor will lose water more rapidly 
by diffusion than it does the more dense sulfur trioxide. When 
concentrated sulfuric acid is cooled to a low temperature, crystals 
of the pure acid, H 2 SO 4 , separate. These melt at 10.5. 


Aqueous Sulfuric Acid. The heat of solution of one mole of 
sulfuric acid in a great deal of water is 75 Kj. So when sulfuric 
acid is mixed with water a great amount of heat is evolved. The 
acid should always be poured into the water, as otherwise so 
much heat is likely to be developed at one point that some of 
the water will be suddenly converted into steam, causing the 
sulfuric acid to spatter. 

When concentrated sulfuric acid is mixed with water consider- 
able contraction occurs so that the total volume is less than 
the sum of the volumes of the water and the acid. The greatest 
contraction takes place when the diluted acid has a composition 
of 70 per cent. A total volume of 100 c.c. before the mixing 
becomes about 97 after the solution of the acid and cooling. 

The vapor pressure of water from aqueous solutions of the 
acid varies continuously with the composition, and is very small 
indeed at ordinary temperatures from the more concentrated acid. 
Upon this fact depends the usefulness of concentrated sulfuric 
acid for drying gases. 

Dissociation of Sulfuric Acid. Pure concentrated sulfuric acid 
is not a good conductor of electricity, but when the acid is dis- 
solved in water it becomes a very good conductor indeed. It 
is a dibasic acid, and like all others of its class, dissociates in 
two ways. In the more concentrated acid, the principal anion 
is HSO 4 ~, but as the acid is diluted this breaks down to give 
H + and S0 4 . As has been pointed out, it is a strong acid, 
but not so strong as hydrochloric. 

Applications of Sulfuric Acid. Enormous quantities of sulfuric 
acid are used in the arts. The largest single use is in the manu- 
facture of fertilizers. A great deal is used in making sodium 
carbonate. In fact, there is scarcely a manufacturing industry 
of any importance which does not directly or indirectly make 
use of sulfuric acid. 

Its many applications depend mainly upon the following facts: 
It is cheap, a fairly strong acid, a moderately strong oxidizing 
agent, not very volatile and forms slightly soluble salts with 
several cations. As an acid it is much used, for example, in the 
preparation of hydrogen, to dissolve the "scale" or coating of 
oxide from metal plates before further treatment in various manu- 


facturing processes and in the preparation of other acids from 
their salts. 

Advantage is taken of its oxidizing power in the separation of 
silver from gold in the process of refining of these metals. The 
silver is oxidized to silver as ion, which then forms silver sulfate, 
while the sulfuric acid is reduced to water and sulfur dioxide. 
The gold is not attacked. 

2Ag + 2H 2 S0 4 = Ag 2 SO 4 + S0 2 + 2H 2 O 

Its extensive use in the manufacture of hydrochloric and 
other acids by the first general method depends upon its slight 
volatility, which more than makes up for the fact that sulfuric 
acid is weaker than many of the acids which it is used to prepare. 
The slight solubility of the sulfates of barium, lead, and calcium 
makes it very useful in the preparation of acids by the second 
general method. In fact, its use in preparing fertilizers depends 
upon the fact that it can convert calcium phosphate into easily 
soluble phosphoric acid or else calcium acid phosphate and 
difficultly soluble calcium sulfate. 

Sulfates. Since sulfuric acid is a dibasic acid, it forms two 
series of salts, acid and neutral sulfates; for example, sodium 
acid sulfate NaHSO 4 and sodium neutral sulfate Na 2 SO 4 . Sul- 
furic acid is a strong enough acid to make the names appropriate 
because the acid salt actually is acid and the neutral salt, neutral. 
The various sulfates will be discussed in some detail under the 
different metals. 

Identification. Barium sulfate is exceedingly slightly soluble 
in water, since it requires more than 300,000 times its weight 
of water for solution. It is promptly formed whenever a solution 
of a barium salt, barium chloride, BaCl 2 , for example, is added 
to a solution of a sulfate or of sulfuric acid. This then is a 
property of the sulfate ion. The barium sulfate is white, and 
is not dissolved in dilute hydrochloric acid, and by this is dis- 
tinguished from all other barium salts, except barium selenate. 


Persulfuric Acid. It will be recalled that when dilute sulfuric 
acid is electrolyzed, hydrogen and oxygen are given off in the 
proportion for the formation of water. 


When 50 per cent, sulfuric acid is electrolyzed, using a very 
small anode and keeping the solution cold, but little oxygen 
appears at the anode. At the same time a new substance, 
persulfuric acid. jIAOfi^ is formed in the solution. The forma- 
tion of this acid is easily understood. Sulfuric acid as concen- 
trated as 50 per cent, or more contains principally the ions H+ 
and HSO 4 ~. Two ionic weights of HSO 4 ~ give up their charges 
to the anode and at once unite to form H 2 S 2 O 8 . If potassium or 
ammonium acid sulfate, KHSO 4 or NH 4 HSO 4 , is used instead of 
the sulfuric acid, the rather difficultly soluble potassium or am- 
monium persulfate soon separates around the anode. These salts 
are now common commercial articles, and find fairly extensive 
use because they are good oxidizing agents. 

Hyposulfurous Acid. The salts of this acid H 2 S 2 4 are formed 
by the action of zinc upon solutions of acid sulfites containing 
an excess of sulfur dioxide, or upon sulfur dioxide in solution 
in absolute alcohol. Sodium hyposulfite, Na 2 S 2 4? is a commer- 
cial article which is extensively used in dyeing. 

The hyposulfites are less powerful reducing agents than zinc, 
but being soluble are much more rapid in their action. 

Thiosulfuric Acid. It will be recalled that a sulfite in solution 
will slowly take up oxygen from the air, passing into a sulfate. 
In much the same way, but more rapidly, a sulfite, say sodium 
sulfite, will take up sulfur forming a thiosulfate, or a sulfate in 
which one atomic weight of sulfur takes the place of an atomic 
weight of oxygen. This is one of the ways in which sulfur shows 
its analogy to oxygen, 

Na 2 S0 3 + S = Na 2 S 2 O 3 

The salt Na 2 S 2 08 -511^0 is known as sodium thiosulfate or incor- 
rectly as sodium hvposulfrfep nr ''fryjW It is the most important 
of the tmosulfates and is much used in photography. It is a 
good reducing agent and is often used as such in place of sulfites. 
For example, it may be used as an antichlor for the removal of 
chlorine from bleached goods, 

' H 2 O + Na 2 S 2 3 + Cla = Na 2 S0 4 + 2HC1 + S 

The protective masks worn by the soldiers in the European 
war to minimize the effect of the " asphyxiating gas/' chlorine, 


were wet with a solution of sodium thiosulfate and carbonate, 
the equation given above represents the main reaction. 

Thiosulfuric acid is so unstable that 'it decomposes almost as 
soon as it is liberated, into sulfur and sulfurous acid. So that 
if a solution of thiosulfate is acidulated it will become milky 
almost at once. This decomposition is largely prevented by 
the addition of sulfites. 

Halogen Compounds of Sulfur. When dry chlorine gas is 
passed over heated sulfur, the two elements combine with the 
formation of sulfur monochloride S 2 Cl2, which is a reddish-yellow 
liquid boiling at 138. It is a very good solvent for sulfur and 
for this reason has been used in vulcanizing rubber. With water, 
it reacts for the formation of sulfur, hydrochloric acid, sulfurous 
acid, and thiosulfurous acid. 

With bromine and iodine, sulfur forms compounds correspond- 
ing to the monochloride. They are the monobromide, SsBi^, and 
thr rnnninHifi8if^:T: Flunrinr and sulfur combine so vigorously 
that they catch fire at ordinary temperatures forming sulfur 

Sulfuryl Chloride. ^Sulfur dioxide and chlorine combine 
directly to form sulfuryl rhloridp T SOgC]^ The combination 
takes place much more rapidly in the sunlight than in diffused 
daylight and is catalyzed by camphor. It is a colorless mobile 
liquid, boiling at 69, and having a density of 1.67. With water 
it reacts to give sulfuric and hydrochloric acids, 

S0 2 C1 2 + 2H 2 O = H 2 S0 4 + 2HC1 

Positive and Negative Valence. Hydrogen and chlorine in 
hydrogen chloride as we have learned are both monovalent. 
When the compound is dissolved in water it dissociates into 
the monovalent hydrogen ion carrying one positive charge, and 
the monovalent chlorine ion carrying a negative charge. This at 
once gives the conception of positive and negative valence, and 
the explanation of the valence of an ion, which then becomes 
identical with the number of positive or negative charges which 
the ion carries. Thus the sulfur ion, S , is negative and is 
divalent, as-is also the sulfate ion, S0 4 . We may, then, say 
that the sulfate ion unites with two atomic weights of hydro- 
gen or sodium as ion; because each of these carries one positive 


charge or is monovalent, and two are required to neutralize the 
charge upon the sulfate ion and to form sulfuric acid or sodium 
sulfate, both of which are electrically neutral. The subject 
of the valence of ions is simple, and this has led to attempts to 
ascribe the phenomena connected with valence in general to 
an electrical origin. In accordance with this view, the elements, 
even in compounds, which are not salts, are supposed to be 
charged, some positively and some negatively and the valence 
of each element is supposed to be determined by the number 
of + or charges which it carries per atomic weight. This 
view seems to be not entirely free from contradictions and may 
be incorrect, but it is in agreement with the facts in the great 
majority of cases and is a very useful conception. According 
to this hypothesis, the oxygen in water and most other compounds 
carries two, the nitrogen in ammonia three, and the carbon in 
methane, CH 4 , four negative charges. 

While it is very easy to decide that the valence of the sulfate 
ion is two, it is not so easy to answer the question "What is the 
valence of the sulfur in this ion?" If we assume that oxygen 
in compounds always has the valence that it has in water, viz., 
two negative, we may. reach a reasonable conclusion in this 
and similar cases. Since the sulfate ion as a whole carries two 
negative charges, it follows that the algebraical sum of the 
positive and negative charges upon the sulfur and the oxygen 
must be 2 or x + a = 2, where "x" represents the charges 
upon the sulfur, and "a" those upon the oxygen. Now from the 
assumption which we have made about the valence of oxygen, 
each atomic weight of this element carries two negative charges, 
and the four would then total eight negative charges, and in the 
equation given above, "a" would have the value 8 

x -8 = -2 

x= +6 

From this it follows that the valence of the sulfur in the sulfate 
ion or in sulfuric acid for that matter, is +6. This is confirmed 
by the fact that sulfuric acid is so simply formed from water 
and sulfur trioxide SOa, in which the sulfur evidently has a va- 
lence of +6, Working in this same way, it is easy to show that 



the chlorine in chloric acid, HC10 3 , has a valence of +5, in 
sodium chlorite, NaC10 2 , of +3, and that the phosphorus 
in phosphoric acid, H 3 P04, has a valence of +5. Nitrogen in 
ammonia NH 3 , and in the ammonium ion NH 4 + , seems to have a 
valence of 3; while in nitric acid HN0 3 , it is +5, so it would 
appear that the valence of an element can not only change in 
value but even in sign. 

The valence of a free element is considered to be zero. 

Oxidation and Reduction. Oxidation and reduction processes 
have been frequently mentioned in what has gone before, but 
have not been discussed in detail. While oxygen is the typical 
oxidizing agent, attention has been called a number of times to 
the fact that we have well marked cases of oxidation into which 
oxygen does not enter in any way. Similarly hydrogen is the 
typical reducing agent, but a great many reduction processes 
are known which do not involve this element. 

A study of these processes soon shows that they alwaysproduce 
changes in the valence of the elements oxidized or reduced. 
That this is so may be easily seen from the table given below. 

An inspection of the last three columns in the first table will 
show that oxidation is always accompanied by an increase in the 
number of positive charges upon the element oxidized, a decrease 
in negative being of course equivalent to an increase in the 
positive charges. The corresponding columns of the second table 
indicate that reduction involves a decrease in the number of 


Valence of the oxidized 














H + 

Zn + + 




KC1O 3 

IO 3 ~ 




Fe + + 

C1 2 

Fe + + + 



+ 1 


MnO 2 



+ 1 

s-- i 2 





Br 2 (H 2 0) 





so 3 -- 



+4 (of S) 

+6 (of S) 




Valence of the reduced 












Cu+ + 


Cu (metal) 




H 2 

Ag (metal) 

+ 1 


I 2 








+5 (N) 

+2 (N) 



Al (in KOH) 

NH 3 




cio 3 - 

so 3 -- 










positive or an increase in the number of negative charges upon 
the element reduced. The most general conception of oxidation 
and reduction then is that oxidation consists in an increase in the 
number of positive charges upon the element oxidized or what 
amounts to the same thing, a decrease in the number of negative 
charges. Reduction consists in a decrease in the number of positive 
charges upon the element reduced, or what amounts to the same 
thing, an increase in the number of negative charges. Oxidation 
and reduction are opposed processes, and one of them cannot 
take place without the other occurring simultaneously and 
to the same extent. This will be seen by examining the above 
table more closely. In each of these reactions, there is an 
oxidizing agent and a reducing agent, and in every case the one 
loses exactly as many charges as the other gains. 


Selenium and tellurium are two elements which bear about 
the same relation to sulfur that bromine and iodine do to chlorine. 
The resemblance of tellurium to sulphur is perhaps less marked 
than that of iodine to chlorine, but selenium furnished fully as 
close a match for sulfur as bromine does for chlorine. A com- 
parison of the atomic weights of the members of these two 
groups of elements is of some interest. 

Chlorine 35 . 46 Bromine 79 . 92 Iodine 126 . 92 

Sulfur. 32.06 Selenium ... 79.20 Tellurium 127.50 

From this it is seen that the atomic weights of sulfur, selenium, 
and tellurium correspond closely to those of chlorine, bromine, 
and iodine respectively. In each group of these elements it be- 
comes easier to liberate an element from its hydrogen compound, 
the greater the atomic weight of the element is. Selenium and 
tellurium correspond to bromine and iodine in that they do not 
occur abundantly in nature although they are rather widely 


Selenium occurs in nature both free and combined. In the 
free state it is found as an impurity in crystals of native sulfur, 
while the compounds are usually associated with the corre- 
sponding sulfur compounds. The element was discovered in the 
flue dust of a sulfuric acid factory, and this together with the 
anode mud of electrolytic lead and copper refineries is still 
the chief source of the element. 

Properties. Like sulfur, selenium takes on a number of 
forms; amorphous, crystalline and metallic. Amorphous sele- 
nium is black and glassy unless finely powdered when it is red. 
There are two kinds of red crystalline selenium; they melt at 



170-180 and are soluble in carbon disulfide. If kept at 150 
for some time, they change into a gray crystalline modification 
melting at 217 and boiling at 690. This gray form has some 
metallic properties and is a poor conductor of electricity in the 
dark, but much better in the light. This is the basis of the sele- 
nium cells which find practical application in automatically turn- 
ing on signal lights at the approach of darkness. Above 1,400 
the molecular weight corresponds to the formula Se 2 ; at lower 
temperature the molecule is more complex. 

Chemical Properties. Selenium combines directly with iron 
and many other metals, and also with hydrogen, oxygen and 
the halogens. As a rule these compounds are formed less 
readily than the corresponding sulfur compounds, and from 
them the element is more easily liberated than is sulfur from 
its compounds. These relations are about the same as those 
between bromine and chlorine. 

Hydrogen Selenide. The methods for the preparation of 
hydrogen selenide, H 2 Se, are very similar to those for hydrogen 
sulfide. It may be formed by the direct union of the elements or 
by acting upon ferrous selenide, FeSe, with hydrochloric acid. 

FeSe + 2HC1 = H 2 Se + FeCl 2 

Hydrogen selenide resembles hydrogen sulfide in that it is a 
colorless poisonous gaseous substance soluble in water, giving 
an acid solution from which the element is deposited upon 
exposure to the air and further that it precipitates difficultly 
soluble selenides from solutions of salts of the heavy metals. 

It is more easily liquefied (boiling-point 41) than hydrogen 
sulfide and has a much worse odor resembling decayed 

Selenium Dioxide and Selenious Acid. Selenium burns in 
air forming solid selenium dioxide, SeC>2, which will dissolve in 
water forming easily soluble crystalline selenious acid. 

It is a weak dibasic acid resembling sulfurous acid. Like the 
latter it acts both as a reducing and an oxidizing agent, but 
differs in that its oxidizing power is much stronger than its 
reducing. It oxidizes sulfurous acid to sulfuric, 

H 2 Se0 3 + 2H 2 S0 3 = 2H 2 S0 4 + Se + H 2 


but reduces permanganate, being oxidized to selenic acid, H 2 Se0 4 , 
4H 2 Se0 3 + 2KMn0 4 = 3H 2 Se0 4 + K 2 Se0 4 + Mn 2 3 + H 2 

Selenic Acid. There is no selenium compound corresponding 
to sulfur trioxide, but selenic acid is strictly analogous to sulfuric 

It is in general a more powerful oxidizing agent than sulfuric 
acid since it liberates chlorine from hydrochloric acid and dis- 
solves gold, but curiously enough, it is without action upon either 
hydrogen sulfide or sulfur dioxide. 


Tellurium is decidedly rare. It occurs free but mostly as 
tellurides of gold, silver, lead and bismuth. Tellurium is more 
metallic in its characteristics than the other members of the 
sulfur family. It forms both amorphous and crystalline modi- 
fications; the latter has a silvery white metallic luster, a density 
of 6.3, melts at 450, and boils at 1,390. At 1,750 the molecular 
weight shows that the formula is Te 2 . The crystalline modi- 
fication is a slight conductor of electricity. The element com- 
bines directly with oxygen, the halogens and many of the metals. 

Hydrogen Telluride. Tellurium compounds are in general 
similar to those of sulfur and selenium, but differ in some im- 
portant points. 

Hydrogen telluride is prepared much as hydrogen sulfide and 
selenide are. It is colorless, gaseous, has a bad odor, boils higher 
(0) than they, and is much less stable. Its solution in water is 
very feebly acid. 

Tellurium Dioxide. Tellurium burns in oxygen forming the 
dioxide which is a white, crystalline substance, slightly soluble 
in water. The solution does not redden litmus. 

Tellurous Acid. Tellurous acid, H 2 Te0 3 , is a white powder, 
slightly soluble in water but dissolved by either potassium hy- 
droxide or strong acids. It then acts both as a base and as an 
acid, but is neither strongly acid nor basic; nothing which acts 
in both these ways is ever strong in either. 

Tellurium trioxide, (Te0 3 ) is known, but unlike sulfur trioxide 
it is totally indifferent toward water. Telluric acid (H 6 TeO 6 
or H 2 Te0 4 -2H 2 O), is a very weak acid with some basic properties. 



It is a powerful oxidizing agent and, unlike selenic acid, attacks 
hydrogen sulfide or selenide. 






S 8 -S 2 

g e8 _Se 2 

Te 2 





Hydrogen compounds. 

H 2 S 


H 2 Se 

H 2 Te 







Very easily 

Most easily 


SO 2 

SeO 2 

TeO 2 





Acid . 

H 2 SO 3 

H 2 SeO 3 

H 2 TeO 3 





Reducing agent 




Oxidizing agent 





SO 3 

TeO 3 


H 2 SO 4 


H 6 TeO 8 

Oxidizing agent 





General. The element nitrogen occurs in nature both free 
and in combination. The free element is found in the air of 
which it constitutes approximately four-fifths by volume. The 
greater part of the remainder is oxygen. The element is de- 
cidedly inactive, but its compounds are among the most reactive 
of chemical substances. This will be seen at once when it is 
mentioned that they include gun cotton, nitroglycerine, and 
almost all of our modern high explosives. The inorganic com- 
pounds of nitrogen which occur in nature are principally am- 
monia, ammonia compounds, and salts of nitric acid. Large 
quantities of very complex carbon compounds of nitrogen are 
present in coal. Nitrogen compounds are of great importance 
to us, since it is with them that the vital phenomena seem to 
be directly connected. The albuminoids constitute an impor- 
tant class of such compounds, and contain almost 15 per cent, 
of nitrogen. 

Preparation of the Element. Impure nitrogen containing the 
members of the argon group can be obtained from air by burning 
phosphorus in the latter^ or a product answering for many com- 
mercial purposes may be secured by the fractional distillation of 
liquid air; but the pure element can be obtained only by decom- 
posing some nitrogen compound. 

Ammonium nitrite, NH 4 N0 2 , when gently heated decomposes 
into water and nitrogen. 

NH 4 N0 2 = N 2 + 2H 2 O 

Ammonia and copper oxide, or nitric oxide, NO, and copper 
react at higher temperatures and yield pure nitrogen. 

Physical Properties. Nitrogen is a colorless, tasteless, odorless 
gas, which is very slightly soluble in water; 100 volumes of the 
latter dissolve 1.6 volumes of nitrogen. The density of nitrogen 
from chemical sources is 0.0012508, while that of the atmospheric 



nitrogen is 0.00125718. It was this difference in density of the 
gas from the two sources that led to the discovery of argon. 
The critical temperature of nitrogen is 146 and the critical 
pressure is 35 atmospheres. Its boiling-point is 195. The 
liquid nitrogen is colorless and has a density, at its boiling-point, 
of 0.8103. At -210, nitrogen freezes to a white solid. 

The molecular weight of nitrogen is 28, and therefore, since its 
atomic weight is 14.01, its formula is N 2 . 

Chemical Properties. Toward most of the other elements, 
nitrogen ordinarily shows itself to be inactive. At higher 
temperatures, or under the influence of the electric glow dis- 
charge it becomes fairly active. When the gas is passed over 
strongly heated lithium, magnesium, calcium, or boron, direct 
combination takes place with the formation of the nitrides of 
these elements, Li 3 N, Mg 3 N 2 , Ca 3 N 2 , and BN. A mixture of 
powdered lime (calcium oxide, CaO) and magnesium reacts so 
vigorously when heated in a current of nitrogen that the pheno- 
mena strongly resemble those of combustion of many substances 
in oxygen. 

When nitrogen is mixed with either oxygen or hydrogen and 
heated, but little action takes place. When an electric spark is 
passed through either of these mixtures, some change occurs, 
but soon comes to an end since the reactions are reversible. In 
the case of nitrogen and oxygen, nitric oxide, NO, is first formed 
and this then combines with the excess of. oxygen present to form 
nitrogen peroxide, N0 2 . The nitrogen and hydrogen unite to 
form ammonia, NH 3 . Nitrogen rendered active by the electric 
glow discharge will continue to glow for some time after the dis- 
charge has ceased, and will react' with phosphorus, sodium, 
mercury, and with some compounds of carbon and hydrogen. 
In many respects it seems to be analogous to ozone, but is less 
stable than the latter. 

Hydrogen Compounds of Nitrogen. By far the most important 
of the compounds of hydrogen and nitrogen is ammonia, NH 3 . 
Some idea of its importance may be gained from the fact that the 
yearly production of the compound and its salts is equal to 
260,000 tons of ammonia. In addition to ammonia we have the 
less important compounds, hydrazine, N 2 H 4 , and hydronitric 
acid, HN 3 . 


Ammonia Occurrence. Ammonia is present in very small 
quantities in the air and in all river and raiii waters, both as 
uncombined ammonia and as ammonium nitrate, NEUNOs, and 
ammonium nitrite, NH 4 NO 2 . 

Preparation. Ammonia may be prepared directly from the 
elements under the influence of an electric spark, or better, by 
heating a highly compressed mixture of nitrogen and hydrogen 
to not over 750 in the presence of finely divided iron or uranium 
as catalyzer. 

An application of the law of mobile equilibrium to the equation 
for the reaction 

N 2 + 3H 2 ^2NH 3 + 50Kj 

shows that since the volume of the ammonia is less than that 
of the elements, high pressure is favorable; while the evolution 
of heat indicates that ammonia will decompose with rising 
temperature, and therefore a low temperature will lead to a more 
complete union of the elements. 

This reaction is now being used for the synthesis of ammonia 
on a large scale. The pressure used is about 200 atmospheres 
and one of the chief problems is to keep the apparatus gas-tight 
at the high pressure and temperature necessary. 

Water acts upon nitrides forming ammonia and the hydroxide 
of the metal. The equation for the decomposition of magnesium 
nitride is given below: 

Mg 3 N 2 + 6H 2 O = 3Mg(OH) 2 + 2NH 3 

A somewhat analogous method consists in the formation of 
calcium cyanamide, CaCN 2 , by the action of nitrogen upon 
calcium carbide, CaC 2 , at fairly high temperatures: 

CaC 2 + N 2 = CaCN 2 + C 

Water decomposes this calcium cyanamide and yields calcium 
carbonate, CaCO 3 , and ammonia, 

CaCN 2 + 3H 2 = CaCO 3 + 2NH 3 

These reactions are the basis of a thoroughly practical process 
for obtaining ammonia from the nitrogen of the air and are now 
being carried out on a very large scale. 





Animal refuse including horns, hides, hair, feathers, etc., 
when heated gives off a part of the nitrogen in the form of am- 
monia. A solution of ammonia in water was formerly prepared 
in this way and was known as spirits of harts- 
horn; a name that is sometimes used at the 
present day. Some ammonia is given off from 
volcanoes and collects around them in the form 
of ammonium chloride, NH 4 C1. 

Coal contains upward of 2 per cent, of nitro- 
gen in combination with carbon, hydrogen, etc.; 
and when the coal is heated in the absence of 
air, as in the manufacture of coal gas, and in 
some of the methods for making coke, a part of 
this nitrogen is given off in the form of ammo- 
nia. The gases are led through water which 
dissolves the greater part of the ammonia and 
also many other substances. This " ammoniacal 
liquor" is then mixed with slaked lime and 
heated. The ammonia is given off and absorbed 
in dilute hydrochloric acid or sulfuric acid, 
forming ammonium chloride, NH 4 C1, or ammo- 
nium sulfate, (NH 4 ) 2 SO 4 . The greater part of 
the ammonia of commerce is prepared from coal 
in this way. 

In the laboratory, .ammonia is usually ob- 
tained by gently heating a mixture of slaked 
lime, Ca(OEn 2 and ammonium chloride (Fig. 
33). Ammonium hydroxide is first formed and 
then decomposed into ammonia and water. 

Ca(OH) 2 + 2NH 4 C1 + CaCl 2 + 2NH 4 OH 
NH 4 OH<=NH 3 + H 2 O 

It may also be secured by heating a strong aqueous solution 
of ammonia. 

Ammonia cannot be dried by either sulfuric acid or calcium 
chloride, since it forms ammonium sulfate with the one and a 
compound, CaCVSNHs, which strongly resembles a salt with 
water of crystallization, with the other. Unslaked lime, CaO, 
or potassium hydroxide is used to dry the ammonia. 

''Ammonium Chloride 
and Slaved Lime 

FIG. 33. 


Physical Properties. Ammonia is a colorless gaseous sub- 
stance whose molecular weight is 17.034. It has a very strong 
characteristic odor which is familiar to everyone. 

Its solubility in water is exceedingly great, one volume at 
room temperature takes up 800 times its volume, or 0.6 of its 
weight of ammonia. At zero degrees, water will dissolve even 
more of the gas. In spite of the great solubility of ammonia, 
it still follows Henry's law approximately, especially at higher 

The critical pressure of ammonia is 111 atmospheres, qdti its 
critical temperature is 132. It may therefore be liquefied by 
pressure alone at ordinary temperatures. The boiling-point is 
33.5. It is a colorless, very mobile liquid which freezes to a 
white, crystalline solid at 77. The density of the liquid at 
its boiling-point is 0.677. The liquid ammonia is a very good 
solvent indeed for many substances, including a large number of 
salts. When salts are dissolved in liquid ammonia, they are 
dissociated very much as they are in water, though to a smaller 
degree at any given concentration than when dissolved in water. 
The ions travel faster in ammonia than in water, so that solutions 
of salts in ammonia are often better conductors of electricity than 
aqueous solutions of the same salts. 

Liquid ammonia is a very common article of commerce, com- 
pressed in strong iron cylinders. The greater part of this is 
used in plants for the manufacture of ice and for cold storage. 
Its use for such purposes depends upon the fact that, like all 
other liquids, it absorbs a great amount of heat when it evapo- 
rates the "heat of vaporization" as it is called. The ammonia 
does not enter into the ice in any way, but is simply used as a 
carrier of heat from a lower to a higher temperature. Such a 
process is not a spontaneous one, and can be brought about only 
by the expenditure of work. The process may be briefly out- 
lined as follows: 

The gaseous ammonia is compressed, generally by a steam 
driven pump (Fig. 34), until it liquefies in a coil of iron pipes 
cooled by a stream of water. In liquefying, it gives out to the 
water the heat of condensation. The liquefied ammonia is then 
allowed to escape into another coil of pipes, called the expansion 
coil, in which it evaporates, thereby taking up from its sur- 



roundings the heat of evaporation. The gaseous ammonia passes 
back from the expansion coil to the pump where it is re-com- 
pressed and liquefied, giving up the heat absorbed in evaporating. 
It then passes once more into the expansion coil and so is kept 
circulating through the system. The expansion coil is usually 
immersed in brine which is cooled by the evaporating ammonia to 
temperatures below the freezing-point of water. Manufactured 
ice is made by placing cans containing about 200 Ib. of distilled 
water in the brine and letting them stand until the water is all 
froz^i The ice prepared in this way is very pure, and much 
superior to most natural ice for domestic purposes. 

* N ! I 

j \ II n i } 1 1 \ hi a H f 


FIG. 34. 

As was mentioned above, the ammonia simply serves to carry 
the heat from a lower to a higher temperature at the expense of 
the work done by the engine. Any other very easily volatile 
liquid might be used instead of the ammonia, and several others 
have been so employed. 

Chemical Properties of Ammonia. When ammonia is heated 
by electric sparks or otherwise, the reverse of the reaction for its 
synthesis (see p. 193) takes place and it is largely decomposed 
into its elements. The decomposition is more complete the 
lower the pressure and the higher the temperature. 

Ammonia will burn in oxygen but not in air. It burns with 
a pale yellow flame forming water, nitrogen, ammonium nitrite, 
NH 4 NO2, and ammonium nitrate, NH 4 NO3. In the presence of 
a suitable catalyzer, platinum for example, ammonia will react 


at a moderate rate with oxygen of the air to form nitric acid. 
This is the basis of what is known as the Ostwald process of 
making nitric acid. 

When ammonia is passed over heated copper oxide or other 
easily reducible oxide, the metal, water, and nitrogen are usually 

When dry ammonia is passed over heated potassium or 
sodium, hydrogen and potassium or sodium amide, KNH 2 or 
NaNH2 are formed: 

2Na + 2NH 3 = 2NaNH 2 + H 2 

The most prominent chemical property of ammonia is its 
ability to combine directly with acids for the formation of 
ammonium salts. For example, with hydrogen chloride, it 
forms dense white fumes of ammonium chloride: 

NH 3 + HC1^NH 4 C1 

With other acids the action is very similar as shown by the 
following equations: 

2NH 3 + H 2 SO 4 <=(NH 4 ) 2 SO 4 
NH 3 + HN0 3 = NH 4 N0 3 

Most of these salts, when heated to a sufficiently high tempera- 
ture, break down into ammonia and the acid from which they 
were formed, just the reverse of the above process. Ammonium 
nitrate and nitrite are exceptions. The decomposition of am- 
monium nitrite into nitrogen and water has already been given. 
The nitrate decomposes into nitrous oxide, N20, and water, 

NH 4 N0 3 = N 2 + 2H 2 

When ammonia dissolves in water, a part of it combines with 
the formation of ammonium hydroxide, NH 4 OH, 

NH 3 + H 2 O^NH 4 OH 

The reaction is reversible, and only a portion of the substances 
are present as hydroxide. At a temperature of 79, ammonium 
hydroxide may be prepared as a white crystalline solid. Another 
crystalline compound of ammonia and water can be obtained at 


a low temperature. It has the composition (NH 4 ) 2 0, and may 
be called ammonium oxide. Its freezing-point is 79. 

Ammonium hydroxide is a weak base and in normal solution 
is dissociated to the extent of about 0.4 per cent. The ions are 
hydroxyl and ammonium, NH 4 + . This ammonium ion has 
many properties which are very similar to those of the potassium 

Upon neutralizing ammonium hydroxide with acid and evapor- 
ating the solution, salts are obtained which are identical with 
those formed by direct union of the acids and ammonia. These 
salts like other salts, are highly dissociated when dissolved in 
water. With solutions of strong bases, they at once react, form- 
ing undissociated ammonium hydroxide which then breaks down 
into the ammonia and water as was shown in the discussion of the 
preparation of ammonia. This reaction is used for the recogni- 
tion of ammonium compounds. The ammonia is detected by its 
action on wet litmus, by the white fumes formed with hydro- 
chloric acid, and even by its odor. 

Ammonium salts, especially the sulfate and chloride, are much 
used as fertilizers to supply the necessary nitrogen for the 
growth of plants. The greater part of the 1,000,000 tons of 
ammonium salts which are annually produced is used as fertilizer. 

An excess of ammonia reacts with chlorine or bromine forming 
ammonium salts and nitrogen. The equation for the reaction 
with chlorine is given below; that for bromine is similar: 

8NH 3 + 3C1 2 = N 2 + 6NH 4 C1 

Hypobromites and hypochlorites act upon ammonia in such a 
way that all the nitrogen is set free in the gaseous state: 

2NH 3 + SNaBrO = 3NaBr + 3H 2 O + N 2 

Hydrazine. Hydrazine, N 2 H 4 , is a colorless liquid which boils 
at 114 and freezes at 1. It combines with water yielding a 
hydrate, N2H 6 O, which is volatile without decomposition. It 
will react with acids with the formation of salts. Two series of 
such salts are known, N 2 H 5 A and N 2 HeAj, in which A stands for 
any-anion carrying one negative charge. Hydrazine is formed 
by the reduction of nitrogen compounds and is itself a very strong 
reducing agent. 


Hydronitric Acid or Hydrazoic Acid. The sodium salt of 
hydronitric acid or hydrazoic acid, as it is often called, may 
be prepared by the action of nitrous oxide, N 2 0, upon sodium 
amide, NaNH 2 : 

NaNH 2 + N 2 O = NaN 3 + H 2 O 

From this salt, the acid is prepared by distillation with dilute 
sulfuric acid, It is a colorless liquid with a strong and very dis- 
agreeable odor. It boils at 37 and explodes violently on contact 
with hot objects. The substance sometimes explodes at ordinary 
temperatures, thus making it very dangerous to handle. 

It is a rather weak acid though somewhat stronger than acetic 
acid. Its salts resemble the chlorides except that the salts of 
the heavy metals are explosive and the difficultly soluble silver 
salt is soluble in the stronger acids, while the chloride is not. 

Air. Air is a mixture of a rather large number of gaseous 
substances and in addition always contains some floating particles 
of dust. The main gaseous constituents which are present in 
almost fixed proportions are nitrogen, oxygen and members 
of the argon group; in addition, it contains variable amounts of 
water vapor, carbon dioxide, and ammonia. In the neighborhood 
of cities and chemical works sulfur dioxide, hydrogen sulfide, 
hydrochloric acid and a few other gases are found. These may 
be regarded as impurities and neglected in the discussion of air. 

The quantity of carbon dioxide varies from about 3 to 4 parts 
per 10,000 in the air of the country to 6 to 7 parts in the cities 
and in badly ventilated rooms may run up to 50 parts. This 
carbon dioxide is formed during the burning or decay of any 
carbonaceous materials, and is also given off during the breathing 
of animals. It is however taken up by green plants, and with 
the aid of the energy absorbed from the light of the sun, is utilized 
in building up the various plant tissues, oxygen being set free 
at the same time. This tends to keep the carbon dioxide content 
of the air approximately constant. 

Composition of the Air. The quantity of water vapor present 
in the air is so variable that no definite statement concerning 
it can be made, but the composition of dry -air is shown in the 
following table: 


Dry Air contains: 

Per cubic meter Per kilogram 

781 . 3 liters nitrogen 755 . 14 grm. 

209 . 9 liters oxygen 231 . 47 grm. 

9.4 liters argon, etc 12.92 grm. 

0.3 liters carbon dioxide 0.46 grm. 

. 1 liters hydrogen ."'. . 01 grm. 

The term argon, etc., includes argon, helium, neon, krypton and 

In spite of the many factors which tend to change the composi- 
tion of the air, the winds produce such a thorough mixing that 
its composition remains practically fixed, although measurable 
variations do occur. 

The density of dry air is 0.001293 under standard conditions. 
In contact with water, each gas dissolves in proportion corre- 
sponding to its solubility and its partial pressure as described by 
Henry's law. The solubility of oxygen is so much greater than 
that of nitrogen that in spite of its smaller partial pressure rela- 
tively more oxygen than nitrogen dissolves. 

Liquid Air. The critical temperature of air is 140 and its 
critical pressure about 39 atmospheres. Liquid air boils from 
194 to 185 according to its composition, and when first pre- 
pared contains from 28 to 50 per cent, of oxygen. As it stands, 
the nitrogen (boiling-point 195) tends to pass off first and 
leaves the oxygen (boiling-point 182.5) behind. This is made 
the basis of very practical methods for obtaining both oxygen 
and nitrogen from the air sufficiently pure for many commercial 
purposes. Liquid nitrogen has a smaller density than that of 
water, while the reverse is the case with liquid oxygen. Liquid 
air when freshly prepared is specifically lighter than water, but 
on partial evaporation becomes heavier. The color also changes 
from practically colorless to the blue of liquid oxygen. 

To liquefy air it is first partially freed from carbon dioxide by 
slaked lime, and then compressed to 150 to 200 atmospheres. 
A great deal of heat is developed in this process so the gas is 
cooled to ordinary temperatures by water in the jackets around 
the compressor (Fig. 35), and by passing it through a coil sur- 
rounded by water. The greater part of the water vapor in the 
air condenses at this high pressure and the remainder together 



with the rest of the carbon dioxide is removed by potassium 
hydroxide. The purified highly compressed air then enters the 
liquefier at the temperature of the room. The liquefier consists 
of a coil of very small copper tubing carrying at its lower end a 
valve for controlling the flow of the air. The whole is inclosed 
in a metal jacket and thoroughly insulated from its surroundings 
by wool. Through the copper tube, the air flows in a continuous 
stream and as it escapes from the valve at the lower end of the 


FIG. 35. 

coil, it becomes somewhat cooler. By the construction of the 
apparatus, this cooled air is compelled to pass up over the coil 
thereby cooling it together with the oncoming air which reaches 
the valve at a lower temperature than the first of the stream, and 
in expanding becomes still cooler and in turn lowers the tempera- 
ture of the coil. This process goes on until a few minutes after 
starting the machine the temperature has been so far lowered 
that about 5 per cent, of the air issuing from the expansion valve 
is liquefied. The rest is returned to the machine, recompressed 
and sent around the cycle once more. In large plants part of the 
cooling is accomplished with carbon dioxide and ethylene. 


The liquid air must not be corked up and will evaporate as 
rapidly as it can get the necessary heat, so after it is obtained it 
is preserved in double-walled vacuum-jacketed vessels which are 
usually coated on the inside of the walls with silver. Such 
vessels are very good heat insulators indeed. 

Liquid air on account of its low boiling-point is an exceedingly 
useful substance for investigations at low temperatures. 

Experiments. Many interesting and striking experiments 
may be performed with liquid air. These depend essentially 
upon two properties; first that it has a very low boiling-point, 
#nd second that it is a source of oxygen. When cooled to the 
temperature of liquid air, most of the familiar gases such as 
chlorine, hydrogen chloride, hydrogen bromide, ammonia, sulfur 
dioxide, methane and hydrogen sulfide become odorless solids. 
Liquids such as alcohol and kerosene freeze, and rubber becomes 
as brittle as glass and may be easily pulverized in a mortar. 
Metals such as copper and steel become a third stronger toward 
a steady pull, but break readily under a quick blow. Crystalline 
metals, on the other hand, are distinctly weaker in every way at 
this low temperature. The heat absorbed during the evapora- 
tion of 1 grm. of liquid air plus the heat required to warm the 
gaseous air to is about 80 cal., so the cooling effect of a pound 
of liquid air in a refrigerator would be about equal to that of a 
pound of ice. 

When liquid air boils, the nitrogen goes off faster than the 
oxygen, and after a time almost pure liquid oxygen is left. 
This furnishes oxygen of such concentration that charcoal or 
steel will burn with exceeding brilliancy in the liquid. Cotton 
or charcoal wet with just the proper amount of oxygen to com- 
plete the combustion will explode when set off with a mercury 
fulminate cap with the violence of an equal weight of dynamite. 
Aluminum powder made into a paste with liquid oxygen will 
burn with a blinding flash upon contact with a flame, and the 
temperature of the mass is changed in a small fraction of a 
second from 183 below to something like 3,000 above zero. 

Air is a Mixture. That air is a homogeneous mixture and not 
a chemical compound is shown by several arguments. 

1. Nitrogen and oxygen may be mixed in the proportion in 
which they are present in the air, and no heat effect will result, 


no changes in volume, nor any other evidence of chemical trans- 
formation and yet the mixture behaves like air in every respect. 

2. The composition of the air varies while that of a compound 
is perfectly fixed. 

3. Relatively more oxygen than nitrogen is dissolved by 
water, while a compound dissolves as a whole, without any 
change in the proportions of the components. 

4. The properties of the air are a mean of the properties of its 
components, which would not be true if it were a compound. 

5. The separation of the gases by liquefaction and fractional 
distillation indicates a mixture. 

Air and Life. Oxygen is absolutely indispensable for the con- 
tinuance of all forms of animal life. The purely aquatic forms get 
their oxygen from air dissolved in the water, while the rest breathe 
the air directly. We, for example, take in about half a liter of 
air at each breath, and remove from this about 5 per cent, of the 
oxygen, giving to it about 3.7 per cent, of carbon dioxide. The 
nitrogen is unchanged. The oxygen taken up in the lungs 
largely enters into a loose combination with a substance called 
hemoglobin in the red corpuscles of the blood and is carried to 
all parts of the body, gradually oxidizing the various body 
substances, chiefly to carbon dioxide, water, and fairly simple 
compounds of carbon, nitrogen, hydrogen, and oxygen. The 
carbon dioxide is largely given up by the blood to the air upon 
its return to the lungs. The oxidation of these substances is the 
source of the body heat, and of the energy which we spend as 

Oxides and Oxyacids of Nitrogen. There are five oxides of 
nitrogen. Three of these are anhydrides of acids. The names 
and formulas of these oxides are given in the following table with 
the names and formulas of the acids opposite their anhydrides. 

Nitrous oxide or hypo- 
nitrous anhydride NO Hyponitrous acid H 2 N 2 O 2 

Nitric oxide NO 

Nitrogen trioxide or 
nitrous anhydride N 2 O 3 Nitrous acid HNO 2 

Nitrogen peroxide NO 2 or N 2 O 4 

Nitrogen pentoxide or 
nitric anhydride N 2 O 6 Nitric acid HNO 3 


Nitric Acid. Nitric acid, the most important of the oxyacids, 
does not occur free in nature, but its salts are present in all 
fertile soils. Potassium nitrate, KN0 3 , has been known for 
centuries under the name of saltpeter. The sodium salt, NaNO 3 , 
occurs in great deposits in the desert regions of Chile, and hence 
is called Chile saltpeter. Enormous quantities of this are 
extracted and shipped all over the world for use in preparing 
other nitrogen compounds, and as a fertilizer. 

Under the influence of electrical discharges nitrogen, oxygen, 
and water will combine for the formation of nitric acid, and this, 
together with the ammonia of the air, is doubtless the source of 
the ammonium nitrate contained in rain water. 

Certain forms of soil bacteria are able to convert ammonia and 
nitrogenous organic substances into nitric acid the latter then 
reacts with the calcium carbonate of the soil to form calcium 
nitrate, Ca(NO 3 ) 2 . 

Nitric acid is prepared by the first general method by distilling 
sodium nitrate with sulfuric acid in cast-iron retorts.. Enough 
sulfuric acid is used to make the sodium acid sulfate which is 
either sold as such or heated with sodium chloride to make hydro- 
chloric^acid and the neutral sulfate. 

NaN0 3 + H 2 SO 4 = NaHSO 4 + HN0 3 

Pure nitric acid is a colorless liquid having a density of 1.53 
at 15. It boils at 86. It gradually decomposes, especially 
under the influence of light into oxygen, water, and oxides of 
nitrogen, which color the acid yellow. The addition of water 
makes the acid much more stable. The conductivity of the pure 
solution is much greater than that of the pure acid which indicates 
that the water has changed the acid into ions. This may account 
for the increased stability. As water is added to the acid, the 
boiling-point gradually rises until when a 68 per cent, acid is 
formed, the boiling-point reaches a maximum at 120. The re- 
lations here are very similar to those between hydrochloric acid 
and water. 

Chemical Properties of Nitric Acid. The chief chemical prop- 
erties of nitric acid depend upon the fact that it is both a strong 
acid and very powerful oxidizing agent. 

As an acid it is as strong as any, being as highly dissociated 


into its ions at any given dilution as hydrochloric acid. Conse- 
quently it is able to do anything that depends upon the concen- 
tration of the hydrogen ions, which any other acid can do. 

As will be recalled, the ordinary reaction between an acid and 
a metal whereby the metal is dissolved and hydrogen is liberated 
is regarded as taking place between the metal and the hydrogen 
ion, and as consisting of the oxidation of the metal to its ion and 
the reduction of the hydrogen ion to the free element. For ex- 
ample, in that between zinc and hydrochloric acid, the chlorine 
ion is considered to take no part, the entire reaction being 
between the zinc and the hydrogen ion, 

Zn + 2H+ + 2C1- <= Zn+ + + 2C1~ + H 2 
or, leaving out, the idle chlorine ion, 

Zn + 2H+<=Zn+ + -{- H 2 

Nitric acid being a strong acid its hydrogen ion is as well able to 
oxidize metals as hydrochloric and in addition, because of its 
being a nitrate, nitric acid is able to oxidize and dissolve copper, 
mercury, and silver which hydrochloric cannot do. It cannot be 
too strongly emphasized that in the case of these last-mentioned 
metals hydrogen is not given off, but in its place nitric oxide ap- 
pears. The nitrogen in nitric acid, in which it has a valence of 
5 positive, has a great tendency to pass into nitric oxide in which 
it has a valence of 2 positive, and in so doing exerts a very power- 
ful oxidizing action, much more so than the hydrogen ion and 
consequently is able to oxidize copper, mercury and silver which 
hydrogen as ion cannot do. The equation for the action on cop- 
per is as follows: 

3Cu + 8HNO 3 = 3Cu(NO 3 ) 2 + 2NO + 4H 2 O. 

In the discussion of the halogens, it was pointed out that con- 
centrated sulfuric acid is a rather powerful oxidizing agent. 
It is much stronger than the hydrogen ion and will oxidize 
copper, mercury, and silver to their ions. As with nitric acid, 
hydrogen gas is not evolved, but an oxide, sulfur dioxide in this 
case, is given off. Here again we have a change in valence, for 
the sulfur in sulfuric acid is 6 positive and in sulfur dioxide 4 


positive. The equation for the reaction between sulfuric acid 
and copper is: 

Cu + 2H 2 S* = CuS0 4 + S0 2 -f 2H 2 

As has been pointed out, a change of valence such as we have 
just mentioned is an essential accompaniment of oxidation. 

Gold and platinum are not oxidized by nitric acid in the 
absence of chlorine as ion, and therefore this acid is used in the 
refining of these metals to dissolve the silver, etc., with which 
they may be alloyed. Because of the vigor of its action on 
metals, nitric acid was formerly called aqua fortis. 

In addition to its use as indicated above, nitric acid is largely 
employed in the laboratory and in the manufacture of nitrates, 
dyestuffs, sulfuric acid, nitroglycerine, guncotton, etc. 

Aqua Regia. Aqua regia is a mixture of nitric and hydro- 
chloric acids, and is capable of attacking gold and platinum, 
metals which do not dissolve in either of these acids alone. The 
action is not due to aqua regia's being primarily a stronger oxi- 
dizing agent than the nitric acid but to the fact that the chlorides 
of gold and platinum are much more stable than their nitrates 
and in aqua regia, the chlorides of the metals may be formed. 
This increased stability of the chloride over the nitrate materially 
assists in the oxidation of the metals and is enough to enable the 
nitrate to do the work, but not sufficient to cause the hydrogen 
ion from hydrochloric acid to oxidize the gold or platinum. 

Aqua regia received its name long ago when the alchemists 
found that it was a solvent for gold which they considered to be 
the king of metals. 

When aqua regia is heated alone, it gives off chlorine and 
compounds of nitrogen, oxygen, and chlorine; one of these is 
nitrosyl chloride, NOC1. 

Nitrates. Nitric acid, as is indicated by its formula, is a mono- 
basic acid and yields only one series of salts, the nitrates. These 
are formed by the action of the acid upon the metals or bases. 

These salts are all readily soluble in water. The nitrates are 
decomposed rather easily at elevated temperatures, generally 
giving oxygen, oxides of nitrogen, and the oxide of the metal 
Most of the applications of nitrates depend upon their action as 
oxidizing agents. When heated with charcoal they often 


detonate. Potassium nitrate, charcoal, and sulfur in the 
proper proportions constitute gunpowder. 

Nitro Compounds. Very concentrated nitric acid acts upon 
compounds of carbon and hydrogen forming nitro-compounds 
and water. For example, nitric acid and benzene, CeHe, react 
for the formation of nitrobenzene, CeHsNC^, and water. 

C 6 H 6 + HN0 3 = C 6 H 5 N0 2 + H 2 

The group N0 2 is known as the nitro group. The water formed 
in the reaction is detrimental, so concentrated sulfuric acid is 
mixed with the nitric acid to take up this water. 

Alcohols (carbon, hydrogen and oxygen compounds containing 
hydroxyl) react with nitric and sulfuric acids for the formation 
of compounds which in their formulas are similar to nitrates, but 
which are not salts. Glyceryl nitrate or nitroglycerine, C 3 H B - 
(N0 3 ) 3 , is formed by the action of this mixture of acids upon an 
alcohol called glycerine, C 3 H 5 (OH)3: 

C 3 H 6 (OH)3 + 3HN0 3 = C 3 H 5 (N0 3 ) 3 + 3H 2 

Nitrogen Pentoxide. Nitrogen pentoxide N 2 5 is a colorless, 
mobile, very volatile liquid which soon passes over into a white 
crystalline solid, melting at 30. It is made by distilling a 
mixture of phosphorus pentoxide and nitric acid. It is the 
anhydride of nitric acid. 

Nitric Oxide. Nitric oxide, NO, is formed by the action of 
somewhat dilute nitric acid, sp. gr. 1.2, upon copper, as shown 
in the following equation : 

3Cu + 8HN0 3 = 3Cu(N0 3 ) 2 + 2NO + 4H 2 O 

The gas so obtained is never pure, since it always contains some 
nitrous oxide, N 2 O, and often nitrogen. 

The pure oxide may be obtained by adding nitric acid to a 
boiling solution of ferrous sulfate, FeS04, in dilute sulfuric acid. 
The equation is as follows: 

6FeSO 4 + 3H 2 S0 4 + 2HN0 3 = 3Fe 2 (S0 4 ) 3 + 2NO + 4H 2 

The ferrous sulfate is oxidized to ferric sulfate, Fe 2 (S0 4 ) 3 , and 
the nitric acid reduced to nitric oxide. 

Properties of Nitric Oxide. Nitric oxide is a colorless gas 
whose molecular wfeight is 30 corresponding to its formula NO. 


It is only slightly soluble in water. The critical temperature of 
the gas is 93, and its boiling-point is 150.2. So it is more 
like the so-called permanent gases, oxygen, hydrogen, etc., than 
most of the gaseous substances which we have been discussing. 

Nitric oxide is the most stable toward heat of the oxides of 
nitrogen. Even when heated to a temperature of 2,000 only 
about 1 per cent, of it is decomposed. In fact it is formed from 
its elements at very high temperatures, and may be obtained by 
passing the electric spark through air. This is made the basis of 
a successful method for the preparation of nitric acid and nitrates 
on a commercial scale as will be explained later. 

Briskly burning wood or phosphorus continues to burn in nitric 
oxide. The flame is about as bright as when the combustion 
takes place in oxygen. Most other burning substances, for ex- 
ample a candle or sulfur, are extinguished. Since the heat of 
formation of nitric oxide is 90 Kj . considerable heat is evolved 
in its decomposition, and a given quantity of a substance like 
phosphorus burning in NO will liberate more heat than if it had 
burned in oxygen. 

Nitric oxide combines directly with oxygen at ordinary tem- 
peratures, forming the reddish-brown nitrogen peroxide, NO 2: 

2NO + 2 = 2N0 2 

This nitrogen peroxide is at once formed whenever the nitric 
oxide comes in contact with the air, consequently it is impossible 
to say anything about the odor of nitric oxide. 

Nitric oxide is soluble in solutions of ferrous salts with the 
production of a dark brown solution containing FeNO ++ . From 
this solution the oxide is driven out by heating to the boiling- 
point. In this way pure NO may easily be obtained from mix- 
tures with other gases, since the latter are either not absorbed in 
the first place or are retained when the solution is boiled. 

The formation of this dark colored solution may be utilized as a 
test for nitric acid, a nitrate, or in general any oxygen compound 
of nitrogen higher than nitric oxide. The equation for the reduc- 
tion of nitric acid to nitric oxide is given on page 207. 

Nitrogen Peroxide. Nitrogen peroxide as has already been 
mentioned is formed by the direct union of oxygen and nitric 
oxide. It may also be prepared by heating certain nitrates. 


Lead nitrate, for example, decomposes to give nitrogen peroxide, 
oxygen, and lead oxide, as shown by the following equation: 

2Pb(N0 3 ) 2 = 4N0 2 + 2 + 2PbO 

It condenses to a yellowish-red liquid which boils at 22 and at 
low temperatures freezes to an almost colorless solid, melting at 

The molecular weight of the gaseous compound varies with the 
conditions of temperature and pressure under which it is de- 
termined. At low temperatures and rather high pressures it is 
nearly 92, corresponding to a formula for the compound of N 2 O 4 . 
As the temperature is raised and the pressure diminished, the 
molecular weight gradually falls until it reaches a limit of 46, 
corresponding to the formula NO 2. This change from one 
form to the other is accompanied by a marked change in color. 
At rather low temperatures the gas has a yellow-brown color, 
as the temperature is rasied it becomes darker and darker 
until at 154 it has reached such a deep black-red that it is almost 
opaque, even in thin layers. 

When nitrogen peroxide is heated above 154, it gradually 
loses its dark color as the temperature rises, and finally becomes 
colorless. This color change is due to the transformation of 
N0 2 into NO and oxygen. 

2N0 2 =* 2NO + 2 

This decomposition becomes practically complete at 620. 

Nitrogen peroxide reacts with water to form nitric acid and NO, 

3N0 2 + H 2 = 2HN0 3 + NO 

In the presence of air, NO will pass into N0 2; which in turn will 
form more nitric acid. So the reaction of NO 2 , air, and water will 
ultimately yield the whole of the nitrogen in the form of nitric 

4N0 2 + O 2 + 2H 2 = 4HNO 3 

Nitric Acid from the Air. When air is passed through an 
electric arc, its temperature becomes extremely high, and a 
small quantity of nitrogen and oxygen unite to form nitric oxide. 
When this has cooled sufficiently, it unites with oxygen to form 




nitrogen peroxide; which then will react with water and oxygen 
as shown in the preceding section to form nitric acid. This 
process has recently become of great importance and is carried 
out at many places, but one of the largest plants is at Notodden, 
Norway. Here the arc is formed by a high potential alternating 
current between copper electrodes which are placed at right 
angles to the poles of a powerful electromagnet. Under these 
conditions the arc is forced to move rapidly alternately up and 
down according to the direction of the current so that it appears 
to form a flat disc of electric flame (Fig. 36) through which the 

Electro Magnei 

FIG. 36. 

air passes. The oxides of nitrogen formed are taken up by 
water in towers similar to the Glover or Gay Lussac towers, and 
the nitric acid is either concentrated and used in this form, or 
is converted into calcium nitrate by letting it act upon slaked 

The calcium nitrate is a good fertilizer and is put on the market 
at a low enough price to compete with the Chile saltpeter. 

Nitrites. When potassium or sodium nitrates are heated, 
especially in the presence of a reducing agent like lead, they lose 
one atomic weight of oxygen per mole and are changed into 

NaN0 3 + Pb = NaN0 2 + PbO 

When an acid is added to a strong solution of a nitrite, the 
nitrous acid formed breaks up immediately into water and nitrous 
anhydride, N 2 3 ; but the latter at once almost completely de- 
composes into nitrogen peroxide and nitric oxide, 
2HN0 2 <= H 2 + N 2 Oa 
N 2 3 * N0 2 + NO 


In this way a nitrite is easily distinguished from a nitrate, since 
the nitrogen peroxide is so highly colored. 

Nitrous acid oxidizes good reducing agents while it is itself 
oxidized by strong oxidizing agents. For example, it oxidizes HI, 
and reduces potassium permanganate, KMnO 4 . The nitrites 
are largely used in the manufacture of organic dyes. 

Nitrous Anhydride. Besides the method just given for the 
formation of nitric oxide and the peroxide in the preparation of 
nitrous anhydride, this same mixture may be obtained by heating 
arsenic trioxide, As20 3 , with nitric acid having a density of 1.30 
to 1.35. The arsenic trioxide is oxidized to the pentoxide, As 2 C>5, 
and the nitric acid reduced to nitrous which at once breaks down 
into water and the mixture of the two oxides. This mixture of 
the two oxides may be condensed by a freezing mixture to a dark- 
blue liquid which boils at 3.5. If the liquid is completely dried, 
it may then be vaporized without dissociating. So dry nitrous 
anhydride, N 2 O 3 , can exist while the moist compound cannot. 

Hyponitrites and Hyponitrous Acid. Sodium hyponitrite, 
Na 2 N 2 O2, may be made by reducing sodium nitrite with sodium 
amalgam. From this hyponitrous acid may be made. It is a 
solid which decomposes irreversibly into water and nitrous oxide 
N 2 O; hence the latter is often called hyponitrous anhydride. 

Nitrous Oxide. Nitrous oxide is readily prepared by heating 
ammonium nitrate, which decomposes into this substance and 

NH 4 NO 3 = N 2 O + 2H 2 O 

It is a colorless gaseous substance which has a faint odor and is 
fairly soluble in cold water. At room temperatures, water will 
dissolve about its own volume of the oxide. Its critical pressure 
is 75 atmospheres and its critical temperature + 36; so it is 
easily liquefied by pressure alone at ordinary temperatures. The 
boiling-point is -90. The molecular weight of the oxide is 
44 and its formula N 2 O. 

Nitrous oxide parts with its oxygen rather more readily than 
nitric oxide so that it will not only continue to support com- 
bustion of brightly burning wood as nitric oxide does, but will even 
cause a glowing splinter to burst into flame. Briskly burning 
phosphorous and sulfur will also continue to burn vigorously in 


nitrous oxide. In all cases, oxides are formed and nitrogen 
set free. 

Nitrous oxide when breathed produces insensibility, and hence 
it is used as an anesthetic. The body does not decompose the 
nitrous oxide in such a way as to utilize the oxygen, so when 
it is to be breathed for any length of time, it is mixed with the 
amount of oxygen required to support life. For use as an 
anesthetic, very pure nitrous oxide is put on the market in 
the liquid form in strong iron cylinders. 

Nitrogen and Life. Compounds of nitrogen are always 
present in all living organisms. They are especially abundant in 
the brain, nerves, and muscles, or, in short, in the tissues which 
seem to be most closely connected with the vital functions. 
Most organisms are entirely unable to make any use of the great 
store of free nitrogen in the air, but must get it from ready 
formed compounds. The plants as a rule get theirs from the 
nitrates, etc., which are present in the soil, while the animals 
obtain the compounds directly or indirectly from the plants. 

Upon the decay of these organisms a portion of the nitrogen is 
liberated in the free state. From this it would seem that the 
supply of available nitrogen compounds must be continually 
decreasing. However, there are two important natural processes 
going on which tend to keep up the supply of nitrogen compounds 
available for plants. The first of these is the formation of 
nitric oxide during the passage of lightning through the air. The 
oxide is then transformed into nitric acid by the action of oxygen 
and water. This combines with the ammonia of the air to form 
ammonium nitrate which is carried by the rain into the soil. The 
second process takes place with the aid of certain forms of bac- 
teria. Some forms are free in the soil, but the most efficient live 
in tubercules on the roots of leguminous plants such as clover, 
peas, alfalfa, etc. These bacteria have the power to transform 
the free nitrogen of the air into compounds which are closely 
related to the albumins, and which may be easily assimilated by 
the plant. Such plants will grow and produce abundant crops, 
and yet leave in the soil an increased amount of nitrogen com- 
pounds. This, then, is the explanation for the well-known fact 
that the fertility of certain soils may be restored by growing upon 
them these leguminous plants. In the older parts of the country, 


it is the custom of farmers to raise a crop of such plants on 
each portion of cultivated land about once in three years. The 
farmers of the* West will doubtless have to adopt this same custom 
in the course of time, if they continue to remove everything from 
the land and return nothing to it, and to burn the stalks and 
straw as so many do at present. 

In addition, as has been mentioned, the compounds of ammonia 
which are obtained in the manufacture of coal gas and coke, as 
well as the calcium cyanamide and calcium nitrate which are 
being made from the nitrogen of the air are available as fertilizers. 
So there seems to be no danger that the earth will lose its fertility 
because of lack of combined nitrogen. 


General. Phosphorus belongs to a group of elements of 
which nitrogen is the member with the smallest atomic weight. 
Jus-t as fluorine, the first member of the halogens, has some 
points of resemblance to chlorine but many of difference, so the 
nitrogen compounds have at least a formal similarity to those 
of phosphorus, but show many marked differences. It might be 
said that the root of the differences lies in the fact that the more 
highly oxidized compounds of nitrogen are unstable and tend to 
go to the less highly oxidized, while just the reverse is the case with 
phosphorus. Phosphorus, however, differs radically from nitro- 
gen in that the element itself is extremely active chemically. 
When exposed to the air even at ordinary temperatures it is 
slowly oxidized, at the same time emitting light. It was on ac- 
count of this property that phosphorus received its name, which 
means light bearer. 

Occurrence. Phosphorus does not occur free in nature, but is 
found widely distributed chiefly in the form of phosphates. The 
most important naturally occurring phosphate is that of calcium, 
Ca 3 (PO 4 )2, which is found in most soils. It constitutes a large 
part of the bones and teeth of animals. Great beds of calcium 
phosphate which have been formed in part at least from the 
fossil bones of animals are found in Florida and Tunis. Enor- 
mous beds, probably the largest in the world, of phosphate rock 
exist in Utah, Montana, Wyoming, and Idaho. A very partial 
survey of the field has located 2,500,000,000 tons of rock averag- 
ing 70 per cent, calcium phosphate. Another mineral which is 
closely related to calcium phosphate is apatite, Ca 5 F(P0 4 ) 3. This 
is a component of many rocks and is found in large quantities in 
Canada. The importance of these compounds of phosphorus 
may be realized when it is known that phosphorus like nitrogen 
is an absolutely essential component of all living organisms, and 




that soluble compounds of this element must be present in a soil 
in order that it may be fertile. In nearly all of the older agri- 
cultural countries, artificial manures containing phosphoric acid 
must be applied to the land. The world mines about 5,000,000 
tons of phosphate rock per year. 

Preparation of Phosphorus. The easiest method of preparing 
phosphorus involves the heating of a mixture of calcium phos- 
phate, sand (silicon dioxide), SiC>2, and carbon to a very high 
temperature in an electric furnace (Fig. 37). Calcium silicate, 
CaSiOs, carbon monoxide, CO, and phosphorus are formed. 

Vapor to 

Calcium Phosphate, 
-' Sand and Carbon 

Calcium Silicate 

FIG. 37. 

The carbon monoxide and phosphorus pass out of the furnace in 
the gaseous state, the phosphorus being condensed under water. 
The calcium silicate is liquid at the temperature of the reaction 
and is drawn off from the furnace from time to time. The whole 
process may be made continuous by feeding irf the reaction mix- 
ture as rapidly as it is used up. The equation is: 

2Ca 3 (PO 4 ) 2 + 6SiO 2 + IOC = 6CaSi0 3 + 10CO + P 4 

The Allotropic Modifications of Phosphorus. The element 
phosphorus exists in two different modifications which are so 
very unlike that it is hard to believe that they are composed of 
the same element. One of these is called ordinary, yellow, 
or sometimes white phosphorus; the other is known as red 
phosphorus. In addition there is a black variety which is much 
like the red. 

Ordinary phosphorus is prepared by the method described 
above and is a slightly yellowish almost colorless wax-like 
solid. It melts at +44 and boils at 270. It is readily soluble 



in carbon disulfide and some other organic solvents, but is nearly 
insoluble in water, under which it is always preserved. From 
its solution in carbon disulfide, crystals of phosphorus may be 
easily obtained. 

When exposed to the air, this form of phosphorus is slowly 
oxidized with the formation of phosphorus trioxide, P 4 6 , and at 
the same time shines so that it may easily be seen in the dark. 
In a limited amount of air, the luminosity will continue until the 
last detectable trace of oxygen has been used up. The tem- 
perature of the phosphorus is but little higher than its sur- 
roundings, so a part of the chemical energy of the change must 
be converted directly into light. The fumes given off have a 

5000 i-mm. Mercury 





100 C 

200 300 

FIG. 38. 


500 C. 

strong garlic-like smell and their continual inhalation pro- 
duces a very serious disease in which the bones of the jaw decay. 
Matchmakers are particularly liable to this trouble. The 
element itself is very poisonous, a tenth of a gram being a fatal 
dose. During the slow oxidation of the phosphorus, a part of the 
oxygen of the air is changed into ozone. If the temperature of 
phosphorus exposed to air is gradually raised, the oxidation 
goes on more and more rapidly until at about 45 it passes into 
rapid combustion. 

The molecular weight of phosphorus, both as vapor and in so- 
lution is 124, and, since its atomic weight is 31.04, the formula 
isP 4 . 


Red phosphorus is very different from the ordinary form. It 
is a dull red crystalline powder and may be obtained by heating 
the yellow variety to 250 in the absence of air. The trans- 
formation is catalyzed by iodine so that in the presence of a 
trace of this substance the transition takes place very rapidly 
indeed at 200. Light also hastens the change, and sticks of 
phosphorus which have been exposed at ordinary temperatures 
to the action of light are covered with a thin layer of red phos- 

Red phosphorus does not melt but passes directly into the 
vapor which is identical with that of yellow phosphorus, and 
yields the latter modification when it is rapidly cooled. In fact, 
it is only by converting the red phosphorus into the vapor and 
condensing the latter that the change from the red into the 
other modification can be made. 

Red phosphorus is not poisonous, is not soluble in carbon disul- 
fide, does not oxidize in. the air at ordinary temperatures, nor 
take fire until heated to about 260. It is, then, very different 
from the ordinary phosphorus, and yet the two modifications can 
be converted the one into the other without any change in weight. 

The yellow modification is at all ordinary temperatures un- 
stable with respect to the red. In accord with this is the fact 
that yellow phosphorus is more soluble than red and has the 
greater vapor pressure as shown by the curves given in Fig. 36. 
The vapor pressure curves do not intersect so there is no transi- 
tion point within the range 100-400, but there is a little evidence 
that a transition point may exist at 77. 

When yellow phosphorus is heated to 200 under a pressure of 
about 4,000 atmospheres it is changed into a more dense variety 
which because of its color is called black phosphorus. This is 
chemically much like the red modification. 

Hydrogen Compounds of Phosphorus. Hydrogen and phos- 
phorus form three compounds, phosphine, PH 3 , liquid hydrogen 
phosphide, P 2 H 4 , and solid hydrogen phosphide, P 4 H 2 . Phos- 
phine is similar to ammonia in its formula and is slightly like it 
in its properties. 

Phosphine may be readily prepared by dropping pieces of cal- 
cium phosphide, Ca 3 P2 (Fig. 39), in water. 

Ca 3 P 2 -f 6H 2 O - 3Ca(OH) 2 + 2PH 3 


The action is very much like that between magnesium nitride 
and water with the formation of ammonia. The addition of a 
little acid facilitates the reaction by rapidly dissolving the 
calcium hydroxide. The phosphine prepared in this way is 
spontaneously inflammable. 

As each bubble of phosphine rises to the surface of the water 
it catches fire with a little explosion and burns to phosphoric acid 

which forms a beautiful smoke ring. 
Pure phosphine is not spontaneously 
inflammable. The compound prepared 
as described above contains small 
quantities of the vapor of the liquid 
hydrogen phosphide, and it is to this 
that it owes the property of catching 
fire on contact with the air. 

Calcium phosphide has found appli- 
cation in connection with life buoys to 
indicate their position at night when 
they are thrown into the water. 

The phosphine may be so far purified 
from the liquid compound by passing it 
through a tube surrounded by a freezing mixture that it will 
not take fire on contact with the air. 

Phosphine is a colorless .gaseous substance which boils at 86, 
smells like decayed fish and is very poisonous. At high tempera- 
tures it is decomposed into its elements. It resembles ammonia 
in its formula and because it combines directly with the hydro- 
halogen acids to form phosphonium salts. It differs from am- 
monia in that it is scarcely at all soluble in water and does not 
form a basic compound with this substance nor does it combine 
with the oxyacids to form salts. 

Phosphonium Compounds. Phosphonium iodide is the most 
stable of these compounds and is easily formed by bringing to- 
gether at ordinary temperatures and pressures phosphine and 
hydrogen iodide, 

PH 3 + HI = PHJ 

It forms beautiful large transparent quadratic crystals which 
sublime at 62. In the vapor state, it is very largely decomposed 


into phosphine and the acid, just as ammonium salts are decom- 
posed into ammonia and the acid. It is a powerful reducing 
agent and finds use as such in the laboratory. 

Phosphonium bromide is very similar to the iodide, and is 
prepared from hydrogen bromide and phosphine at ordinary 
temperatures and pressures. It sublimes at 38. 

Phosphonium chloride cannot be prepared by bringing hydro- 
chloric acid and phosphine together at ordinary temperatures 
and pressures. At 35 it can be obtained at ordinary pressures, 
but requires a total pressure of at least 18 atmospheres at 14 to 
produce it or to keep it from decomposing into its components. 
Phosphonium compounds are decomposed by water into phos- 
phine and the acid. From this we gather that the phosphonium 
ion, PH4 + , is not stable. 

Liquid Hydrogen Phosphide. Liquid hydrogen phosphide 
boils at 57. The vapor is spontaneously inflammable. Its 
analysis and molecular weight indicate that the formula is P2H 4 . 

Halogen Compounds of Phosphorus. Each of the halogens 
will combine directly with phosphorus forming fairly stable 

Chlorine, bromine and iodine form trihalides, PC1 3 for example, 
which are all liquids. Fluorine, chlorine and bromine form penta 
compounds, PC1 5 for instance; with the exception of the fluoride 
which is gaseous, these are solids. 

Each of these compounds is decomposed by water with the 
formation of an oxyacid of phosphorus and the hydroacid of the 
halogen. It will be recalled that this property was made use of 
in the preparation of hydriodic and hydrobromic acids. The 
chlorine compounds are rather more important than the others 
and will be described in some detail. 

Phosphorus trichloride, PC1 3 , is formed by passing chlorine 
over phosphorus contained in a retort to which is connected a 
receiver. The combination takes place with the evolution of 
so much heat that the greater part of the trichloride formed 
distills over and condenses to a colorless liquid in ths receiver. 
It boils at 76 and has a density of 1.6. It reacts with water and 
other compounds containing oxygen and hydrogen to form phos- 
phorous acid and hydrogen chloride. Because of this it fumes 
when in contact with moist air. 


Phosphorus pentachloride, PC1 5 , is prepared by the action of 
an excess of chlorine upon the trichloride. It is a pale yellowish- 
green solid which when heated under atmospheric pressure does 
not melt, but passes directly into the vapor without going through 
the liquid state. This is due to the fact that the temperature at 
which the vapor pressure of the solid is equal to the pressure of 
the air is lower than its melting-point. We might say that its 
boiling-point is lower than its freezing-point. When the vapors 
are cooled they pass directly back to the solid state. That kind 
of distillation in which the solid passes directly into the vapor and 
reappears in the solid state upon cooling without the liquid 
modification coming in as an intermediate step is called sublima- 
tion. By heating the pentachloride under presssure, the boiling- 
point may be so raised that it lies above the melting-point which 
is 166. In the vapor state, phosphorus pentachloride, PC1 6 , 
is partially decomposed into chlorine and the trichloride. 

Phosphorus pentachloride fumes strongly in contact with air 
and has a very irritating action on the mucous membrane. It 
rapidly reacts with water, forming phosphoric acid, H 3 P0 4 , and 
hydrogen chloride. 

Oxygen Compounds of Phosphorus. The oxides of phosphorus 
are the tetroxide, ?204, and trioxide, P20a, or P^e, and the 
pentoxide, PzOs or P 4 Oi . The simpler formulas for the last two 
oxides are the ones which are the more commonly used for these 
substances, but the more complicated ones correspond to their 
molecular weights. 

Phosphorus pentoxide is a white powder which is formed 
when phosphorus burns in dry air. It is exceedingly hygro- 
scopic, and instantly combines with water to form metaphos- 
phoric acid, HPOs. It is the most thorough drying agent known, 
much better than sulfuric acid and far better then calcium 
chloride. Owing to its great affinity for water, it will even with- 
draw the elements of water from compounds containing hydrogen 
and oxygen. Its use in the preparation of nitric anhydride 
depends upon this property. 

The trioxide is formed when phosphorus burns in a limited 
supply of air. It differs from the pentoxide in that it has a low 
melting-point, 22.5, and boiling-point, 173, and hence may 
be easily separated from the higher oxide by distillation. It is 


colorless. With cold water it reacts very slowly forming phos- 
phorous acid, and hence is often called .phosphorous anhydride. 

Acids of Phosphorus. Phosphorus pentoxide is the anhydride 
of three important acids, metaphosphoric acid, HPO 3 , pyro- 
phosphoric acid, H 4 P 2 7 , and orthophosphoric acid, H 3 PO 4 . 
These are formed by the interaction of varying quantities of 
water and the anhydride: 

P 2 O 5 + H 2 = 2HP0 3 
P 2 Os + 2H 2 = H 4 P 2 O 7 
P 2 O 5 + 3H 2 = 2H 3 P0 4 

In addition to these three acids there are the less important ones, 
phosphorous acids, H 3 P0 3 , hypophosphoric acid, H 4 P 2 O6, and 
hypophosphorous acid, H 3 PO 2 . The pentoxide is not the anhy- 
dride of these last three acids and it will be noticed that the 
valence of the phosphorus is 3, 4 and 1 respectively. 

Orthophosphoric acid is of greater importance than any other 
acid of phosphorus and is always in mind when phosphoric acid 
is mentioned without qualification. It may be prepared as in- 
dicated above, or obtained from its salts by the second general 
method for the preparation of acids. The first general method 
is not available since the acid cannot be distilled. It is most 
conveniently made in pure state by oxidizing ordinary phos- 
phorus with dilute nitric acid, nitric oxide being formed at the 
same time. By concentrating the solution, the excess of nitric 
acid is driven off. The orthophosphoric acid obtained in this 
way is a thick viscous liquid which crystallizes slowly and with 
difficulty. The melting-point of the pure acid is 41. 

Its solution in water reddens litmus and has a pure agreeable 
acid taste. It is not very strongly dissociated. A solution 
of one gram equivalent of phosphoric acid in a liter of water 
contains only about one-fourteenth as much hydrogen as ion as 
a hydrochloric acid solution of the same concentration. 

Orthophosphoric acid is a tribasic acid as is indicated by its 
formula, H 3 PO 4 . It forms three series of salts according as one, 
two, or all three of the atomic weights of hydrogen per mole are 
replaced by metals. We have then normal salts and two kinds 
of acid salts. These latter are distinguished by using Greek 
numerals to show how many atomic weights of metal are present. 


Thus we have normal sodium phosphate, NasPO^ which is often 
called trisodium phosphate; disodium phosphate, Na 2 HPO 4 ; and 
monosodium phosphate, NaH 2 PO4. The names tertiary, second- 
ary and primary are also used to designate these salts. Solutions 
of the mono salt are slightly acid in reaction while those of the 
disodium salt are faintly alkaline, and of the tri- exceedingly so. 
These phenomena are connected with the fact that phosphoric 
acid is a weak acid and also polybasic. Like all other polybasic 
acids, it dissociates one atomic weight of hydrogen more strongly 
than the rest. The dissociation H 3 PO 4 = H+ -f- H 2 PO 4 ~ takes 
place to a fair degree, while that of H 2 PO 4 ~ = H + + HPO 4 is 
slight, and HPO 4 scarcely breaks down at all into H + and 
P0 4 . Because of the weakly acid character of the ions 
HPO 4 and H 2 P0 4 ~, the trisodium salt is almost completely 
hydrolyzed into the disodium salt and sodium hydroxide, and 
the disodium phosphate is also slightly affected. 

The phosphates which occur in nature are the .normal salts 
and the principal one is that of calcium, Ca 3 (P0 4 )2. This salt 
is almost universally present in the soil. It is, however, very 
slightly soluble so that it is only slowly available as a source of 
phosphorus compounds for plants. Monocalcium phosphate, 
Ca(H 2 PO 4 )2, is easily soluble and makes an excellent fertilizer. 
It is known in commerce as "superphosphate" and is prepared 
by treating calcium phosphate with dilute sulfuric acid from the 
lead chambers. So extensive is the use of phosphate fertilizers 
that the larger part of the sulfuric acid made is used in their 

Pyrophosphoric Acid. When orthophosphoric acid is carefully 
heated to about 250, it loses water and is changed into pyro- 
phosphoric acid, H 4 P 2 0?, 

2H 3 P0 4 = H 4 P 2 O 7 + H 2 

When dissolved in water, it gradually reacts with the latter form- 
ing the ortho acid. The formula indicates that it is tetrabasic, 
and it forms four series of salts, NaH 3 P 2 07, Na 2 H 2 P 2 07, Na 3 - 
HP 2 O 7 , and Na 4 P 2 0?. Of these the second and fourth are the 
best known. 

Metaphosphoric Acid. Metaphosphoric acid, HPO 3 , may be 
obtained by heating orthophosphoric acid to a higher tempera- 


ture than that required for the preparation of pyrophosphoric 
acid, or by adding the proper amount of water to the pentoxide. 

Metaphosphoric acid melts at a rather high temperature and 
solidifies to a wax-like mass on cooling. It is put on the market 
in the form of transparent sticks and is known in commerce as 
glacial phosphoric acid. These sticks are hard and glass-like 
because they contain a little sodium metaphosphate. At high 
temperatures it passes into vapor, and in this state has the for- 
mula (HPO 3 ) 2 as is shown by its molecular weight. 

At room temperatures, a solution of metaphosphoric acid is 
changed rather rapidly into pyrophosphoric acid and the latter 
then more slowly into the ortho acid. 

The sodium salt of metaphosphoric acid is formed by heating 
monosodium orthophosphate to a high temperature, 

NaH 2 P0 4 = NaPO 3 + H 2 O 
or by heating sodium ammonium phosphate, NaNH 4 HPC>4, 

NaNH 4 HPO 4 = NaPO 3 + NH 3 + H 2 

Sodium metaphosphate fuses to a clear glass-like mass and in the 
fused state is able to dissolve most of the oxides of the metals, 
receiving from them in many cases, colors characteristic of the 
metals. This property is made use of in blowpipe analysis. 

Phosphorous Acid. Phosphorous acid, H 3 P0 3 , is formed by 
the action of cold water upon phosphorus trioxide, P 4 Oe, or of 
water upon the tri-chloride, bromide, or iodide of phosphorus. Its 
formula indicates that it is a tribasic acid, but it will form only 
two series of salts, NaH 2 PO 3 and Na 2 HPO 3 . This property 
is probably due to an extreme case of the difference in the degree 
of dissociation of the hydrogens of a poly basic acid, to which 
attention was called under phosphoric acid. 

Phosphorous acid is a crystalline solid melting at 70. When 
heated strongly, it decomposes into phosphine, metaphosphoric 
acid, and water. It is a powerful reducing agent, being oxidized 
to phosphoric acid. 

Phosphorus Sulfides. By heating red phosphorus and sulfur 
together in the proper proportions in the absence of air, a rapid 
reaction takes place, and according to the relative proportions of 
the substances, the sulfides P 4 S 3 , P 2 S 3 , P 3 S 6 , and P 2 S 5 are formed. 


Of these sulfides, P 4 S 3 is the most important, because of its use 
in matches. It does not take fire spontaneously, but does so 
when gently heated. Red phosphorus is used because the reac- 
tion is too violent with yellow. 

With water the sul fides react to form hydrogen sulfide and the 
acids of phosphorus. 

Applications of Phosphorus. By far the greater part of the 
phosphorus made is used in the manufacture of matches. These 
may be divided into two classes, those which will strike any- 
where, and those which will ignite only when rubbed on a specially 
prepared surface. The first kind are made by dipping the sticks 
into melted paraffin, and then into a paste containing glue, 
powdered glass, iron oxide, lead oxide, and phosphorus sesqui- 
sulfide, P 4 S 3 . The phosphorus sulfide is enclosed within the mass 
and so protected from the air. Upon striking the match the 
friction raises the temperature, at the point of contact, above the 
kindling point of the phosphorus sulfide. The whole very com- 
bustible mixture in the head then burns and sets fire to the wood. 
These matches are so easily ignited that they are liable to produce 
accidental fires. This type of match was formerly made with 
free phosphorus and was highly poisonous, but the kind contain- 
ing phosphorus sulfide is practically non-poisonous. 

The second kind is the safety match which lights on the box. 
The head of this match may contain potassium chlorate and 
dichromate, sulfur, manganese dioxide, iron oxide, powdered 
glass, glue, and gum arabic, but no phosphorus. The striking 
surface on the box is composed of a mixture of red phosphorus, 
antimony trisulfide, dextrine, and lampblack. Safety matches 
may sometimes be ignited by friction on a non-conducting surface 
such as that of glass. Since these matches contain no ordinary 
phosphorus they are relatively non-poisonous and can be manu- 
factured without injury to the workmen. For the reason that 
it is practically impossible that they should be ignited accident- 
ally, this form of match is much superior to the other for domestic 
purposes. In the laboratory red phosphorus is used wherever 
possible, because it is more moderate in its action and less danger- 
ous than the yellow. 


General. Carbon belongs to a family of elements, which in 
their highest state of oxidation are tetravalent; the other mem- 
bers of the family are silicon, germanium, tin, and lead. Carbon 
and silicon are non-metallic elements, while the others are metals 
and will be discussed later. The compounds of carbon are 
exceedingly varied, more than a hundred and fifty thousand of 
them being known. In fact, the chemistry of carbon is so com- 
plex that it is considered a separate branch, and is called organic 
chemistry. It received this name because all living organisms 
contain carbon compounds, and it was for a long tune supposed 
that most of these substances could be produced only by the 
life processes of such organisms. At the present day so many of 
these compounds have been sjoithesized in the laboratory that 
we are now of the opinion that it is possible to prepare any carbon 
compound artificially. 

The great importance of the compounds of carbon may be 
realized when we consider that nearly all of our foods, and most 
of our medicines, consist essential^ of these substances, and 
that the principal source of the energy which we use in our 
various industries is the carbon in coal, gas, oil, etc. 

Occurrence. Carbon occurs free in nature in two distinct 
allotropic crystalline modifications, which are known as diamond 
and graphite. A third form of elementary carbon, known as 
amorphous carbon, or charcoal, may be obtained by heating 
organic substances, such as wood, sugar, etc., to a high tempera- 
ture in the absence of air. Many persons consider that amor- 
phous carbon is present in coal and so would say that this form 
of carbon occurs in nature. In combination, carbon is found as 
an essential constituent of all living organisms; as carbon dioxide 
in the air and water; as calcium carbonate in the great beds of 
limestone; as very complex compounds chiefly in combination 
15 225 


with hydrogen in coal and oil; and as methane, CH 4 , in natural 

Diamond. As is well known, diamond is a clear colorless 
substance which crystallizes in octahedrons and is exceedingly 
hard. Because of these properties, its rarity and the fact that it 
has a very brilliant luster, diamond has long been valued highly 
as a gem. The principal localities in which diamonds are found 
at the present day are Brazil and South Africa. In Brazil, they 
are obtained from alluvial deposits, while in South Africa they 
occur within and in the immediate vicinity of old volcanic 
chimneys. The diamonds have apparently been formed from 
carbon under great pressure, at the high temperature of the lava. 
They are usually small, very rarely weighing more than a few 
grams; but one, the Cullinan diamond, was found in 1905 at 
Premier mine in Transvaal which weighed 620 grm., or 3,024 
carats, or 1.37 lb., and even then the stone was evidently merely 
a fragment of a much larger one. 

Very small diamonds have been prepared artificially by 
Moissan. He placed charcoal made from sugar in cast-iron 
heated to a very high temperature in an electric furnace and then 
cooled the iron as rapidly as possible. Under these conditions a 
rigid shell of solid cast-iron was quickly formed around the out- 
side of the mass, while the interior still remained molten. Now 
cast-iron upon solidification increases in volume, so that, as the 
interior gradually froze, great pressure was produced. Under 
these conditions the carbon was simultaneously exposed to very 
high temperature and pressure, and a small portion of it crystal- 
lized as minute diamonds. Because of its extreme hardness it is 
used for cutting and writing on glass and many other purposes 
where such a property is required. The density of diamond, 
3.5-3.6, is greater than that of the other forms of carbon. When 
diamonds are heated in the air, or oxygen, they burn to carbon 
dioxide. Toward most chemical reagents it is very indifferent. 

Graphite. Graphite is the second crystalline modification of 
carbon. It occurs in many localities, chiefly Austria and Ceylon, 
in rather poorly formed rhombohedral crystals. It is black in 
color, non-transparent, so soft that it may be readily scratched 
with the finger nail, and leaves a black streak on paper, has a 
density of 2.25, and so differs radically in its properties from 



diamond. It is prepared artificially on a large scale by heating 
petroleum coke, anthracite coal, or any other nearly pure form 
of carbon, to a very high temperature in the electric furnace (Fig. 
40) . Graphite is highly resistant to chemical action, but rather 
less so than diamond. When heated in the air, or oxygen, it slowly 
burns to carbon dioxide. When heated alone it undergoes no 
change until the highest temperature of the electric arc is reached, 
when it passes slowly into vapor without melting. Because of its 
infusibility, graphite when mixed with a small portion of clay to 
act as a binder, is used in making crucibles which have to stand 

.'Molded Carbon Electrodes to be changed -to Graphite 

FIG. 40. 

high temperatures. Because of these same properties and the 
further fact that it is a good conductor of electricity, graphite 
is extensively used at present for making electrodes for use in 
electric furnaces and in the electrolytic industries. Crystals 
of graphite very easily break up into fine scales which slip over 
one another with but little friction. On this account it is often 
used as a lubricant. Finely ground graphite mixed with clay 
and slightly baked, constitutes the "lead" of our lead pencils. 
The more clay the harder the pencil. 

Amorphous Carbon. As has been mentioned above, amor- 
phous carbon is prepared by heating certain carbon compounds 
to a high temperature in the absence of air. A rather large 
number of different forms are distinguished according to their 
state of division and the substances from which they were formed. 
We have, for example, charcoal of various kinds, soot or lamp 
black, and coke. The density, hardness and electrical conduc- 
tivity of all forms of amorphous carbon increase with the tem- 
perature to which they are heated. At the highest temperatures 
of the electric furnace, the various forms of amorphous carbon 
are transformed into graphite. Like graphite, amorphous carbon 


is infusible. It is, however, much more easily combustible than 

Soot is a comparatively pure form of finely divided carbon and 
is formed by the combustion of various carbon and hydrogen 
compounds with an insufficient supply of air. It is very soft, 
intensely black, has a small density, is easily combustible, and is 
a good non-conductor of heat or electricity. It finds application 
as a pigment in printer's ink, paint, etc., under the name of 

Charcoal. The most common form of charcoal is that pre- 
pared by heating wood to fairly a high temperature in cast iron 
retorts. In addition to the charcoal, methyl or wood alcohol, 
CH 3 OH, acetic acid, CH 3 COOH, and creosote are formed and are 
condensed, separated and sold for more than the value of the 
charcoal. Bone charcoal or bone-black as it is called, is made 
by carbonizing bones in retorts. At the same time a valuable 
by-product known as bone oil is produced. All solid substances 
seem to have a tendency to increase, right at their surface, the 
concentration of any gaseous substance with which they may be in 
contact. This property is called adsorption. It is shown 
strongly by all forms of charcoal, especially by those such as bone 
and cocoanut shell charcoal which have per unit weight very 
large surfaces owing to the fineness of the pores of the materials 
from which they were made. The higher the molecular weight 
and boiling point of the gas and the lower the temperature, the 
greater is the adsorption. 

When cooled to the temperature of liquid air, cocoanut shell 
charcoal will adsorb very completely all gases except hydrogen, 
helium and neon. One of the most convenient ways for produc- 
ing a high vacuum is to seal a bulb containing cocoanut charcoal to 
the vessel to be exhausted and then cool the bulb with liquid air. 
In a short time, the air in the vessel will be so completely taken 
up by the charcoal that a very high vacuum will be produced. 

When charcoal is added to solutions it adsorbs nearly every 
substance present. This is more pronounced with bone-black 
than other charcoals for the reason given above. Here again, 
the adsorption is more complete the higher the molecular weight 
of the substance affected, and since coloring matters are usually 
very complex, they are adsorbed to a greater extent than most 


other substances. Crude sugar contains a coloring matter which 
is removed in refining by boiling a solution of the sugar with 
bone-charcoal. Some of the sugar is adsorbed, but not enough 
to make a serious loss. Charcoal is a catalyzer for many reac- 
tions, particularly between gases. Here again we are dealing 
with a surface phenomenon, and the charcoals prepared from 
the more finely cellular substances are the more active. 

Coal may be considered natural charcoal, and has probably 
been formed by the partial decomposition of vegetable matter 
under water. Various varieties of coal are distinguished, such 
as anthracite, semi-anthracite, bituminous, and lignite. These 
differ rather widely in the relative proportions of carbon, hydro- 
gen, nitrogen, oxygen, sulfur, and mineral matters which they 
contain, the anthracite being the nearest pure carbon, and the 
others decreasing in purity in the order named. The world's pro- 
duction of coal is about 1,300 million tons and that of the United 
States about 500 million tons each year. 

Coke is formed by heating bituminous coal with the exclusion 
of air until practically all the volatile matter has been distilled 
off. The gases given off are rich in compounds of carbon and 
hydrogen, and are very valuable for fuel and illuminating pur- 
poses. In addition, they contain many useful products, such 
as ammonia, benzene, C 6 H 6 , coal tar, etc. Coke is largely used 
as a fuel in blast furnaces and many other metallurgical opera- 
tions, since it produces a very high temperature, burns without 
smoke and contains a relatively small amount of sulfur. 

Relation Between the Allotropic Modifications. Although the 
various modifications of carbon differ so radically in their prop- 
erties they are really composed of the one element carbon, and 
contain nothing else, as is shown by the following argument. 
When equal weights of pure diamond, graphite, or amorphous 
carbon are burned in oxygen, equal weights of carbon dioxide are 
produced in each case, and nothing besides carbon dioxide is 
formed in any case. The heat of combustion of the various 
modifications is different so that in this, as in other cases, the 
allotropic modifications differ in their energy content. Because 
of the non-volatility and insolubility of carbon, it is impossible 
to determine its molecular weight. Its atomic weight, however, 
is 12.005. 


Oxygen Compounds. Carbon forms two well known oxides, 
the monoxide CO, and the dioxide C02. The second of these 
compounds is by far the more important. 

Carbon Dioxide. Carbon dioxide, CO 2 , is a gaseous substance 
whose molecular weight is 44. It is colorless, has a very feeble 
taste and odor, its critical temperature is 31, and its critical 
pressure is 73 atmospheres. 

Liquid carbon dioxide is a common article of commerce. It is 
sold in very strong cylinders. At a temperature of 20 its vapor 
pressure is 58.5 atm. or 860 Ib. to the square inch, so it may be 
easily seen that the cylinder should be well constructed and 
handled with care. The principal use for this carbon dioxide is 
in the manufacture of soda water and other carbonated bever- 
ages. The boiling-point, or rather subliming-point of carbon di- 
oxide is 78.2; the freezing-point is considerably higher than 
this 57; so that at ordinary pressures liquid carbon dioxide 
cannot exist in open vessels. Liquid carbon dioxide is a colorless, 
very mobile substance, which is but slightly soluble in water. 
It is not a conductor of electricity, does not act as a solvent 
toward most substances, and does not redden litmus paper; in 
short, it does not present any striking chemical properties. 

Solid carbon dioxide may be formed by allowing the liquid 
modification to flow out from the cylinder at any tempera- 
ture below the critical temperature of the liquid. As soon as 
the liquid comes in contact with the air, the greater portion of it 
is at once turned into vapor. The heat necessary for this 
transformation is absorbed largely from the carbon dioxide 
itself, and so far cools a portion of this that it assumes the solid 
state, although its temperature in this condition is 78.2. 

The solid carbon dioxide is a white snow-like substance which 
is very convenient for use as a cooling agent, since we can so 
easily get with it the definite temperature of 78.2. In order 
to secure better thermal contact with the substance to be cooled, 
it is commonly mixed with ether. By diminishing the pressure 
upon the carbon dioxide, temperatures lower than 78.2 may 
be obtained. In this way by causing it to evaporate in a vacuum, 
a temperature of 100 may be reached. 

Carbon dioxide occurs in nature in very large quantities. The 
air contains three to four parts of this substance to ten thousand 


parts by volume. It escapes in great volumes from volcanoes, 
mineral wells and springs, and at various places on the earth it 
issues in an almost pure state from fissures. It is formed during 
the breathing of animals and in the fermentation and decay of 
organic substances. It may be prepared by burning carbon or 
any carbonaceous material in an excess of oxygen or of air; or 
by heating limestone, calcium carbonate, CaCO 3 , in lime kilns 
where it breaks down into calcium oxide, CaO, and carbon diox- 
ide, CO 2 ; 

CaCO 3 ^ CaO + CO 2 

or by the action of an acid upon a carbonate. The car- 
bon dioxide which is put on the market in the liquid form is very 
largely obtained from lime kilns and from the fermentation 
vats of the breweries. The ordinary laboratory method for its 
preparation is to place some pieces of marble, which is a very 
pure form of calcium carbonate, in a flask and allow hydro- 
chloric acid to run slowly upon them. Under these conditions 
the carbon dioxide is given off in a steady stream, 

CaC0 3 + 2HC1 = CaCl 2 + CO 2 + H 2 O 

Being heavier than air and not very soluble in water, it may be 
collected either by the displacement of air or over water. Carbon 
dioxide is neither a supporter of combustion nor of animal life. 
Air containing 4 per cent, of carbon dioxide will extinguish a 
candle flame, but will support respiration for a short time. Since 
carbon dioxide is heavier than air and is being gradually given 
off by the earth, it often accumulates in old wells, cellars, cisterns, 
etc., in quantities sufficient to render it dangerous to persons 
entering such places. The presence of carbon dioxide in danger- 
ous quantities may be easily detected by the lowering of a lighted 
lantern. If the flame is extinguished no one should enter until 
the excavation has been thoroughly ventilated. 

It dissolves in its own volume of water at ordinary tempera- 
tures and the solubility follows Henry's law. 

The solution has a feebly acid reaction and contains the very 
unstable substance, carbonic acid, H 2 C0 3 . This is a very weak 
dibasic acid which forms two series of salts, the normal and 
the acid carbonates. Because of the weakness of carbonic acid, 
the easily soluble normal salts are rather strongly hydrolyzed and 


are alkaline in reaction; and the so-called acid salts are almost 
neutral in their reaction. The most abundant naturally occur- 
ring carbonate is that of calcium, CaCO 3 . Sodium carbonate, 
Na 2 CO 3 , is manufactured in enormous quantities from sodium 
chloride, and is a very important chemical. Sodium, potassium 
and ammonium carbonates are soluble in water. Most of the 
other carbonates are not soluble." 

When an acid is poured on a carbonate, the carbonic acid 
formed decomposes at once with a brisk effervescence due to the 
escape of carbon dioxide. Such an effervescence suggests a carbon- 
ate but does not prove its presence because a few other salts, 
'sulfites and sulfides, for example, will also effervesce. The best 
test for carbon dioxide is the formation of a white precipitate 
of calcium carbonate when it is brought in contact with lime 
water, which is a solution of calcium hydroxide in water. The 
carbon dioxide is first dissolved by the water forming carbonic 
acid, which then reacts with the calcium hydroxide for the for- 
mation of water and the difficultly soluble calcium carbonate 

CO 2 + H 2 O <=> H 2 CO 3 
Ca(OH) 2 + H 2 CO 3 = CaCO 3 + H,O 

Because of the fact that carbon dioxide will combine with water 
for the formation of carbonic acid, it is often called carbonic 

Circulation of Carbon. During the breathing of animals 
oxygen is taken up from the air in the lungs forming an un- 
stable compound with the hemoglobin of the blood. It is then 
carried to all parts of the body and gradually oxidizes the com- 
plex compounds contained therein, forming carbon dioxide as 
one of the products; this is carried by the blood to the lungs 
and there given up to the air and expired. In the dark 
practically this same process takes place in plants, but in the 
light those plants which contain green coloring matter in their 
leaves are able to bring about the reverse change, partially 
breaking up the carbon dioxide absorbed from the air, liberating 
a part of its oxygen in the free state and combining the residue 
with water for the formation of complex compounds of which 
the sugars, starches, and celluloses are prominent examples. 
These compounds are food for plants as well as for all animals, 


which live directly or indirectly upon them. This decomposition 
of carbon dioxide and the building up of complex compounds from 
it takes place only under the simultaneous influence of chloro- 
phyl, the green coloring matter of plants, and of light. It is 
accompanied by the absorption of a great amount of energy 
which is obtained from the light of the sun. The energy so 
absorbed is given out once more when these complex compounds 
burn, decay, or oxidize in the bodies of animals; and so the ulti- 
mate source of the energy of animals and of all engines, which 
are driven by the burning of wood and other vegetable products, 
is the light of the sun. Coal has been formed by the partial 
decomposition of vegetable organisms, and the energy obtained 
from it was originally absorbed by these organisms from the 
light of the sun ; so the steam engine of to-day is really driven by 
the energy of the sunlight of thousands of years ago. 

By this cycle of changes, carbon dioxide is successively taken 
up from and given back to the air, and thus there is in nature a 
sort of circulation of carbon. The energy, however, accompany- 
ing these transformations apparently does not move in any such 
cycle; that given out during the decomposition of the complex 
compounds is not available for the re-formation of these sub- 
stances, and this change can take place only with the influx of 
a fresh supply of energy from the sun. We are to think of the 
energy relationships in the following way: A stream of energy 
comes to the earth in the sunlight; a very small part of this is 
absorbed by the green plants and stored as chemical energy in 
the form of complex carbon compounds, which are available for 
use as foods, or combustibles, and so furnish us with the greater 
part of the energy which we employ The other sources of 
energy which man employs, water power and the wind, are 
really dependent upon the sun, so that we owe practically all 
of our stores of available energy to this body. 

Photochemical Action. Red and yellow light rays are prac- 
tically the only ones which have any influence on the reactions 
which take place in plants; while in the photochemical changes in 
photographic plate, the red rays are almost inactive and blue, 
green and violet produce most of the effect. From these and 
similar instances, it follows that every kind of light can produce 
some kind of chemical change. 


Carbon Monoxide. Carbon monoxide is a gaseous substance 
which is formed when coal is burned in a limited supply of air. 
It may also be prepared by the action of carbon at a high tem- 
perature upon carbon dioxide. The carbon monoxide burns in 
the air with a pale blue flame, and it is the burning of this sub- 
stance which produces the blue flames that are so commonly seen 
at the top of a hard coal fire. The chemical changes taking place 
in such a fire are mainly as follows : When the air first comes in 
contact with a hot carbon at the bottom of a firebox, carbon 
dioxide is formed, which, passing up through the body of the coal, 
is reduced to carbon monoxide; this burns when it comes in 
contact with the air at the surface of the fire. The equations for 
the reactions are as follows, 

C + O 2 = CO 2 and CO 2 + C=>2CO 

These reactions take place on a large scale in the manufacture 
of " producer gas" for fuel and power purposes, see p. 250, and 
also in the iron blast furnace. 

When steam is passed over highly heated carbon, carbon mon- 
oxide and hydrogen are produced as shown in the equation 

C + H 2 O = CO + H 2 

This is one of the reactions which takes place in the technically 
important process for making what is known as water gas. This 
will be discussed in some detail later, see p. 250. 

Carbon is an excellent reducing agent and is much employed 
for this purpose in the extraction of metals from their oxides. 
When this change takes place at a high temperature, as is 
usually the case, carbon monoxide is formed. The equation for 
the reduction of zinc oxide is 

ZnO + C = Zn + CO 

Carbon monoxide is often prepared by heating oxalic acid, a 
white crystalline substance, whose formula is H 2 C 2 04, with 
sulfuric acid. Under these conditions the oxalic acid decomposes 
into carbon dioxide, carbon monoxide and water, 

H 2 C 2 O 4 = CO 2 + CO + H 2 O 

The carbon dioxide is removed by passing the mixture of gases 
through a strong solution of sodium hydroxide. Formic acid 


heated with sulfuric acid breaks down into water and carbon 

HCOOH = CO + H 2 

In this reaction and that of oxalic acid the sulfuric acid appar- 
ently does nothing except take up the water formed. 

Properties. Carbon monoxide is a colorless, tasteless gas with 
a very faint, peculiar odor, very slightly soluble in water. Its 
molecular weight 28, critical temperature 140 boiling point 
190 are nearly the same as those for nitrogen. Carbon monoxide 
will combine directly with oxygen, chlorine, and many other sub- 
stances. With oxygen, carbon dioxide is formed with the 
evolution of much heat, making it an excellent fuel. It combines 
with chlorine forming the substance, COC1 2 , carbonyl chloride 
or phosgene. This combination takes place when carbon monox- 
ide and chlorine are exposed to sunlight, or when a mixture of 
the two gases is passed over animal charcoal which acts as a 

Carbon monoxide is a powerful reducing agent. The reduc- 
tion of iron oxide to metallic iron by this gas takes place on an 
enormous scale in the blast furnace. The reaction may be 
represented by the following equation: 

Fe 2 3 + SCO = 2Fe + 3CO 2 

Carbon monoxide is extremely poisonous. This is connected 
with the fact that it will combine with the hemoglobin of the 
blood to form a very stable compound which is unable to take 
up the oxygen of the air. So poisonous is carbon monoxide that 
700 c.c. of the gaseous substance will be sufficient to kill an 
average man. 

Carbon Disulfide. When charcoal is heated to a rather high 
temperature in the vapor of sulfur, the elements combine forming 
an easily volatile liquid substance, called carbon disulfide, CS 2 . 
This substance is now manufactured on a large scale in electric 
furnaces (Fig. 41). It is a colorless liquid whose boiling-point 
is 46. It has a very high index of refraction for light. When 
perfectly pure it has a pleasant ethereal odor; the commercial 
article, however, contains impurities which gives it a very 
disagreeable smell. It is an excellent solvent for sulfur, iodine, 
phosphorus, rubber, resins, and fats and hence finds extensive use 



in the arts and manufactures. A mixture of carbon disulfide 
vapor and air is violently explosive and is remarkable for the fact 
that ignition takes place at temperatures as low as 260. On 
this account even greater care must be taken in handling this 
substance than is used with gasoline. 

Nitrogen Compounds of Carbon. At the very high tempera- 
ture of the electric arc, carbon and nitrogen will combine directly 

with the formation of a gaseous com- 
pound, cyanogen, C 2 N 2 . This sub- 
stance may also be obtained by heating 
mercuric cyanide, Hg(NC) 2 . This 
breaks down into metallic mercury and 
cyanogen, as shown in the equation, 

Hg(NC) 2 = Hg + C 2 N 2 


Cyanogen is a colorless gas, having 
a peculiar pungent odor resembling 
that of peach kernels; it is exceedingly 
poisonous and burns in the air with a 
characteristic red-violet flame, forming 
carbon dioxide and nitrogen. Its 
stability at high temperature is due to 
its being formed from the elements 
W1 ^ the absorption of heat. 

Chemically, cyanogen resembles the 
halogens. It forms a series of salts 
called cyanides, which are similar to the 

chlorides, but contain the group NC, instead of chlorine. Silver 
cyanide, for example, is a difficultly soluble white substance, which 
is very much like silver chloride in its general properties. From 
these salts, hydrocyanic acid or prussic acid, HNC, may be easily 
obtained by the addition of a less volatile acid and distilling. 
It is a colorless liquid which boils at 27 and has a very strong 
odor like that of peach pits. It is one of the weakest acids. 
On this account, the soluble cyanides react strongly alkaline 
when in solution in water, owing to hydrolysis. Hydrocyanic 
acid and the cyanides are extremely poisonous substances. 
The poisonous effect of the peach pits is due to hydrocyanic 
acid formed by the action of water upon a substance called 


FIG. 41. 


amygdalin which is in the pit. Not only is hydrocyanic acid 
a poison toward animals, but it will also stop or "poison" the 
action of many catalyzers, and it is thought that perhaps the 
poisonous action of this substance on animals is due to its stop- 
ping the catalytic processes which are going on within them. 

Cyanates. Cyanides are good reducing agents, and when 
oxidized often form compounds called cyanates, potassium 
cyanate, KNCO, for example, is formed by cautiously oxidizing 
potassium cyanide, KNC. Cyanic acid, corresponding to these 
cyanates is very unstable. It is a colorless liquid with a strong 
smell resembling that of acetic acid. 

Thyiocyanates. When a cyanide is heated with sulfur, the two 
substances combine to form a thiocyanate. These correspond 
to the cyanates. The formation of potassium thiocyanate, 
KNCS, is shown in the following equation, 


The thiocyanates are used as reagents for detecting the ferric 
salts with which they give an intense red coloration. 

The Hydrocarbons. Several hundred compounds of carbon 
and hydrogen have been described. For purposes of classifica- 
tion and as an aid in getting a comprehensive view of their rela- 
tions, they have been divided into several series according to the 
composition and general properties of the substances. The 
principal series is that known as the Marsh Gas or Paraffin series, 
and the simplest member of this is marsh gas or methane. This 
series is distinguished by the fact that in all the compounds 
belonging to it carbon has a valence of 4 the highest valence 
which it is capable of taking on. For this reason, the members 
of this series are sometimes called the saturated hydrocarbons. 
Figure 42 gives the names and formulas, and indicates the melt- 
ing- and boiling-points of some of the more important of the lower 
members of the normal paraffins. 

An inspection of this figure will show that the boiling- and 
melting-points of these hydrocarbons increase as they become 
more complex. It will be noticed, too, that the formula for each 
succeeding member of the series differs from that of the one which 
just precedes it by CH 2 . A series of carbon compounds which 
are related to one another in this way is called a homologous 



series. The members of such a series have similar properties 
which gradually change with the composition of the compound. 
The change in the boiling-points of the compounds shown in 
the figure is an illustration of this. Most of these hydrocar- 
bons occur in nature in large quantities, the majority of them 
being found in crude petroleum, while methane and ethane are 
present in natural gas. 

Pentadecane C^Ifx 


TetradecaneC u /4 


Tridecane C u ff K 



Oodecane C u H K 


Undecane C n Hu 



Decane C IO I& 



Nonane C 9 H^ 





Octane C s H u 

Heptane C 7 H,,. 



Hexane C 6 H lt 


Pentane C 5 fl[ 2 





Butane C 4 H lo 



Propane C 3 H s 


Ethane C, H, 



Methane C Jf 4 



-200 -100 100 200 300C. 

Fia. 42. 

Crude Petroleum. Crude petroleum or "oil" is found in 
many places in the United States, in Ontario, in Baku on the 
Caspian Sea, in Roumania, in India, and in Japan. This oil is a 
mixture of a great many substances, but principally of hydrocar- 
bons, and usually consists very largely of the hydrocarbons of the 
marsh gas series. In oil refining, for commercial purposes, no 
attempt is made to prepare pure substances from the oil, but the 
more volatile portions are distilled off and separated into frac- 
tions according to their density and adaptability for particular 



purposes. The following table gives the name, approximate 
components, density, and uses for some of the more common 


Trade name 

Chief constituent 






Solvent) automobiles. 




Solvent; automobiles. 



722-0 737 

Solvent; automobiles 



745-0 824 

Traction and station- 


Nonane, decane, etc. . . 


ary engines. 
Illuminating ; fuel ; and 
stationary engines. 

The residue left after removing these substances is then 
distilled in a current of steam. Some of the higher boiling 
products which come over under these conditions are used 
for lubricating oils, but must first be purified from the solid 
members of the series by cooling to a low temperature and 
filtering with the aid of filter presses. The solid substances so 
secured constitute what is known as paraffin, while the low 
melting product, which is salve-like at ordinary temperatures, is 
called vaseline. The final residue left in the retort is used as 
asphalt. Natural asphalt is a mixture of hydrocarbons which 
occurs in nature. Just how these hydrocarbons are formed 
in nature is not known, but they have probably been produced, 
in most cases at any rate, by the partial decomposition of animal 
and vegetable organisms, although it is possible that they may 
have resulted from the action of water upon carbides, compounds 
of carbon with metals. 

Most of the hydrocarbons are rather indifferent chemically. 
They are not acid like hydrogen chloride nor basic like ammonia. 
They of course burn in oxygen, forming carbon dioxide and water, 
and react with chlorine or bromine to form compounds. When 
heated to a very high temperature they all show a tendency to 
break down into hydrogen, carbon and simpler hydrocarbons 
containing a relatively smaller amount of carbon. The process 


is technically known as " cracking" and is now used on a large 
scale to increase the yield of gasoline. 

Methane. Methane or marsh gas, CH 4 , is the first member of 
the paraffin series. It is the principal constituent of natural 
gas. It is formed during the decay of vegetable matter in the 
bottom of ponds and marshes, and to this fact it owes its name, 
marsh gas. It is found in considerable quantities in coal, and 
since its mixture with air is explosive, it is one of the causes 
of the disastrous coal mine explosions. When found in mines 
it is often called "fire-damp" by the miners, while the carbon 
dioxide, which is formed when it burns and is left in the mine 
after an explosion, is called "choke-damp." The critical tem- 
perature of methane is -82 and its critical pressure 56 atmos- 
pheres; its boiling-point is -164. Because of its very low 
critical temperature, the methane of natural gas does not exist 
in the depths of the earth in a liquid condition, but is present 
simply as a highly compressed gas, filling the pores of the rocks. 

The origin of methane found in nature is as much in doubt as 
that of the other hydrocarbons, but the probabilites are that it 
was formed by the decomposition of organic matter. It may be 
made by the action of water upon aluminum carbide, 

A1 4 C 3 + 12H 2 = 4A1(OH) 3 + 3CH 4 

and it is possible that a part of the methane has been produced 
by this reaction. In the laboratory the gas is usually prepared 
by distilling a dry mixture of sodium acetate and sodium hy- 
droxide. The equation for the reaction is as follows, 

NaC 2 H 3 O 2 + NaOH = Na 2 C0 3 + CH 4 

Methane burns in the air forming carbon dioxide and water. 
The flame has very little luminosity, unless used with a Wels- 
bach mantle composed of the oxides of cerium and thorium. 
Because of the great amount of heat which is produced during 
the combustion, methane is very valuable for fuel and in the 
form of natural gas is much used for this purpose. As will be 
seen by the following equation, CH 4 + 2O 2 = C0 2 + 2H 2 O 
one volume of methane will require two volumes of oxygen or 
ten volumes of air for complete combustion. The explosive 


qualities of a mixture of gas and air are utilized as a source of 
energy in the gas engine. 

When a mixture of chlorine and methane is exposed to sun- 
light the four combining weights of hydrogen are successively 
replaced by the chlorine forming the following series of com- 

Methyl chloride CH 3 C1 

Methylene chloride CH 2 C1 2 

Chloroform CHC1 3 

Carbontetrachloride CCU 

The other product of each of these reactions is hydrogen chloride. 
These are not ionized when dissolved in water and hence are not 
salts. Chloroform is perhaps the most important of these 
substances and is generally prepared by the action of bleaching 
powder upon ordinary alcohol. It is a colorless, easily volatile 
liquid which is used as a solvent and as an anesthetic. Carbon 
tetrachloride is now being manufactured on a large scale, but not 
directly from methane. It is a valuable solvent and finds 
many applications as such, being used as a substitute for gaso- 
line in dry cleaning, because it has the great advantage that it is 
not inflammable. In fact it is now being used as a fire ex- 
tinguisher and is about the only thing beside carbon dioxide 
which will put out a gasoline fire. 

Radicals. A study of the derivatives of methane reveals the 
existence of a series of compounds which have the common 
characteristic that their formulas can be so written that each 
contains a CH 3 group, CH 3 C1, CH 3 Br, CH 3 I, CH 3 OH, (CH 3 ) 2 O. 
When these compounds undergo certain transformations the 
CH 3 group seems to pass unaltered from compound to com- 
pound and so acts like a chemical individual, being like the 
ammonium ion in this respect. A group which behaves in this 
way is called a radical, and this particular group, CH 3 , is known 
as the methyl radical. A study of the chemistry of the carbon 
compounds shows that we have to do with a large number of 
such radicals. For example, each member of the paraffin series 
will form a monovalent radical, which contains one less atomic 
weight of hydrogen per mole than the hydrocarbon from which 
it was formed. Ethane, C 2 H 6 , forms the ethyl radical, C 2 H 6 , 



propane, C 3 H 8 , yields the propyl radical, C 3 H 7 , etc. These 
radicals, which we have been considering are not ions, and have 
the remarkable property of being able to combine with almost 
any element or radical irrespective of the electrical character of 
the substance. For example, the ethyl radical will combine 
with sodium to form sodium ethyl, C 2 H 5 Na, or with chlorine 
forming ethyl chloride, C 2 H 5 C1, and yet sodium maybe considered 
a typical electropositive element, while chlorine seems to be 
decidedly electornegative, at least in the chlorides. This 
tendency toward universal combination exhibited by these 
radicals is apparently one of the causes for the great variety 
of carbon compounds. The methyl radical, for example, will 
combine with chlorine, bromine, or iodine, forming methyl 
chloride, bromide, or iodide, with the hydroxyl, producing 
methyl alcohol, CH 3 OH, and with oxygen yielding methyl ether 
(CH 3 ) 2 O. If no other substance is present for it to combine 
with, one methyl radical will combine with another to form the 
compound CH 3 CH 3 or C 2 H 6 , which is ethane, and this in turn 
will yield the ethyl radical, C 2 H 5 , which under the proper condi- 
tions will combine with a methyl radical to form the compound, 
C 2 H 5 CH 3 , or C 3 H 8 , which is propane, the third member of the 
group. Proceeding in this way, practically the whole series 
of paraffin hydrocarbons has been built up and consequently 
they all may be regarded as derivatives of methane. 

Other Hydrocarbons. Among the hundreds of hydrocarbons, 
mention may be made of benzene, CeH 6 , and toluene, C6H 5 CH 3 . 
These are obtained by the distillation of coal and are of great 
importance because they are the starting-points for the prepa 
ration of a great many very useful substances. From benzene 
we may make carbolic acid, CeH 5 OH, aniline, CeHsNH^, and 
the aniline dyes. Toluene yields among other compounds, 
trinitrotoluene, CeH 2 (NO 2 ) 3 ,CH 3 , which is a powerful explosive. 
Types of Carbon Compounds. The principal carbon com- 
pounds with which we shall have to deal may be divided into a 
few well-marked types; the hydrocarbons, which we have already 
discussed, alcohols, ethers, aldehydes, acids, esters and carbo- 
hydrates. The following table gives the general formulas for 
these compounds. In this table " R" stands for any monovalent 
radical, and "A" for the anion of a monobasic acid. 



General formulas 




R 2 O 









The Alcohols. The first member of the alcohol group, methyl 
or wood alcohol, CH 3 OH, is one of the by-products obtained in 
the destructive distillation of wood for charcoal. It is a color- 
less liquid, which boils at 66. It is valuable as a fuel and as a 
solvent. It is decidedly poisonous and is largely used in the 
crude state to mix with ethyl alcohol for the production of the 
so-called denatured alcohol. 

The molecular weight and analysis of methyl alcohol show that 
it contains per molecule one atom each of carbon and oxygen 
and four of hydrogen. Experience teaches that one of the 
hydrogens is different from the others, because sodium will 
replace only one, forming the methylate, CHsONa, and further 
that hydrochloric acid will react with the alcohol, yielding water 
and methyl chloride, CH 3 C1. To represent the action with 
sodium, one of the hydrogens must be written differently from 
the others, say HCH 3 0; but the action with the acid suggests a 


hydroxide, and the formula, CH 3 OH or H C OH, indicates 

both actions, and the latter is called the structural formula for 
methyl alcohol. In a similar way, the properties of other com- 
pounds are studied, and structural formulas devised to represent 

Ethyl alcohol, C2H 5 OH, or grain alcohol as it is called, is 
formed by the action of an enzyme or catalyzer produced by 
yeast upon sugars, C 6 Hi 2 O 6 . The process is called fermentation 
and the equation for the reaction is as follows: 

C 6 H 12 O 6 = 2C 2 H 5 OH + 2CO 2 


The alcohol so formed may be largely separated from the water 
by distillation, the alcohol passing over first since its boiling-point 
is 78.3. The alcohol of commerce contains about 95 per cent, 
by volume of alcohol and 5 per cent, of water. It is not possible 
to separate these two substances completely by distillation, since 
they form a minimum boiling mixture, boiling at 78.15 and 
containing 4.43 per cent, of water. (See hydrogen chloride 
and water.) Absolute alcohol free from water may be obtained 
by adding unslaked lime to the 95 per cent, alcohol. The lime 
combines with the water to form calcium hydroxide, and after the 
action is complete, the absolute alcohol may be distilled off. The 
density of alcohol at 25 is 0.7851. Alcohol is a valuable fuel and 
an extremely useful solvent for many substances which are diffi- 
cultly soluble in water. It is largely employed for the prepara- 
tion of medicines. Denatured alcohol is 95 per cent, alcohol, 
mixed with some substance which cannot be readily separated 
from the alcohol and which will render it totally unfit for use as a 
beverage. A very commonly used formula is 89.5 per cent, 
grain alcohol, 10 per cent, wood alcohol and 0.5 per cent, benzine. 
Such alcohol is exempt from the payment of any internal revenue 
tax, and is consequently much cheaper than the pure article, and 
so finds extensive application in the arts. 

The Ethers. The general formula for these substances is R 2 0. 
Ethyl ether, (C 2 H 5 )2O, is the best known and by far the most 
important member of this group. It is prepared by heating a 
mixture of ethyl alcohol and sulfuric acid to 140. The equation is 

2C 2 H 5 OH = (C 2 H 5 ) 2 O + H 2 O 

The sulfuric acid acts as a catalyzer and will change many times 
its own weight of alcohol into ether. Ethyl ether is a colorless 
liquid boiling at 34.6. It is used as solvent for fats, etc., and as 
an anesthetic. 

The Aldehydes. The general formula for the aldehydes is 
RCOH. These substances, in general, may be formed by the 
partial oxidation of the alcohols. The first member of the series 
is formaldehyde, HCOH. In this case hydrogen takes the place 
of the "R" in the general formula. Formaldehyde is produced 
by the incomplete combustion of methyl alcohol when the vapors 


of this substance mixed with air come in contact with platinum 
gauze or platinized asbestos. Formaldehyde is gaseous at ordi- 
nary temperatures, the boiling-point being 21; it has a very 
pungent, irritating odor, is quite soluble in water and is put 
upon the market in a 40 per cent, solution, known as formalin. 
Both the solution and the gaseous substance are powerful germi- 
cides and are used as disinfectants. 

The Acids. The general formula for the organic acids is 
RCOOH. These substances are numerous and important. 
The first member of the series is formic acid, HCOOH. As 
was the case with formaldehyde, in the formic acid the hydro- 
gen takes the place of "R" in the general formula. Sodium 
formate, HCOONa, is made by passing carbon monoxide over 
heated sodium hydroxide, the equation being 


From this salt, the acid may be obtained by distilling it with 
dilute sulfuric acid; concentrated sulfuric acid decomposes 
formic acid into water and carbon monoxide. It is a colorless 
liquid boiling at 100, freezing at 8.6. Like almost all other 
carbon acids it is a weak acid, although somewhat stronger 
than acetic acid which is the next member of the series. In 
spite of the fact that, as its formula indicates, it contains two 
atomic weights of hydrogen per mole, it is a monobasic acid, 
only one of the hydrogens being ionizable. This fact that the 
hydrogens differ in their properties is expressed by writing them 
in different positions in the formula HCOOH instead of COOH 2 . 
As indicated by the general formula for the series, the organic 
acids contain the group COOH, and to this is ascribed their acid 
character, and the hydrogen of this group is supposed to be that 
which ionizes. 

Acetic Acid. Acetic acid, CH 3 COOH, is the second member. 
It is formed during the destructive distillation of wood, and is 
one of the valuable by-products of this process. Large quan- 
tities of the substance are produced by the oxidation of ethyl alco- 
hol by the oxygen of the air with the aid of a certain form of 
bacteria, known as B. Aceti. The common name for this is 
mother of vinegar. The dilute solution of acetic acid, formed 
by the oxidation of the alcohol contained in cider, is known as 


vinegar. Vinegar usually contains 4 to 5 per cent, acetic acid. 
Pure acetic acid is a colorless liquid which boils at 119 and 
freezes at 16.7. It mixes with water in all proportions. As 
has been mentioned repeatedly acetic acid is a weak acid, being 
only slightly dissociated. Although it contains four atomic 
weights of hydrogen per mole, it is a monbasic acid, three of the 
four atomic weights of hydrogen being present in the methyl 
radical, and these are non-ionizable. 

The Esters. When an alcohol is acted upon by an acid, a 
reaction takes place which consists in the formation of water, 
apparently from a combination of the acidic hydrogen of the acid 
with the hydroxyl of the alcohol, and of an ester a compound 
produced by the union of the radicals from the acid and the 
alcohol. The equation for the reaction between ethyl alcohol 
and acetic acid is as follows: 

CH 3 COOH + C 2 H 5 OH & H 2 O + CH 3 COOC 2 H 5 

is called ethyl acetate, and is an important mem- 
ber of the group of esters. Many of these are pleasant smelling 
substances and are largely responsible for the odors and flavors 
of our flowers and fruits. In addition the fats and vegetable 
oils are esters of the alcohol glycerine, so one of our most impor- 
tant classes of foods is composed of esters. 

Carbohydrates. The carbohydrates are an extensive group 
of carbon, hydrogen and oxygen compounds in which, with very 
few exceptions, the hydrogen and oxygen are present in the same 
relative proportions as in water, so that the general formula 
for nearly all members of the group may be written C m H 2n O n . 
The great importance of this class of compounds will be recog- 
nized at once when it is known that it includes sujcrose (cane or 
beet sugar) Ci 2 H 22 On, glucose, C 6 Hi 2 O 6 , starch, (CeHioOs)*, and 
cellulose, (CeHioOs)!,. Cotton and linen are practically pure 
cellulose, and the woody part of all plants consists largely of this 
substance. In wood it is associated with another compound 
called lignin of unknown composition. If the wood is treated 
with a solution of sodium or calcium acid sulfites the lignin is 
dissolved and the cellulose is left behind in fine fibers suitable for 
making wood pulp paper. Starch is obtained from many plants, 


chiefly corn, wheat, potatoes and rice. Its importance as a food 
and for laundry work is well known. 

When boiled with a dilute acid starch takes up water and 
changes to glucose; a thick syrup of the latter made from corn 
starch is often called corn syrup. Cane and beet sugar are iden- 
tical; their source and general properties are too well known to 
require discussion beyond the fact that when boiled with dilute 
acid, water is taken up and a solution of glucose and fructose is 
formed. This process is called inversion of sugar and the product 
is known as invert sugar. 

Some Explosives. The importance of explosives for industrial 
and war-like purposes is universally recognized. Besides the old- 
fashioned gunpowder (p. 300), and nitroglycerine (p. 207), there 
may be mentioned the following: Guncotton, [C^H^O^NOs^]*, 
is made by acting on cotton with a mixture of nitric and sulfuric 
acids. It is very explosive and is used in making smokeless 
powders; one form of these is known as walsrode. It is made of 
a mixture of 99 per cent, guncotton, and 1 per cent, ethyl acetate. 
Picric acid, C 6 H 2 (NO 2 ) 3 OH, is made by acting on phenol, i.e., 
carbolic acid, (C 6 H 5 OH), with nitric and sulfuric acids. It is the 
explosive known as lyddite. Trinitrotoluene, C 6 H 2 (N0 2 ) 3 CH 3 , 
is made in the same way using toluene, CeHsCHs, in place of 
phenol. It is known as T.N.T. It may be melted and poured 
into shells. When it explodes, great volumes of black smoke 
are produced. Dynamite is nitroglycerine absorbed in some por- 
ous material such as finely divided silica, kieselguhr as it is called. 
Nitric acid is used in the preparation of each of these explosives, 
and this is one reason for its great industrial importance. 

Unsaturated Hydrocarbons. The methyl radical may be 
imagined as being formed from methane by the dropping of one 
atomic weight of hydrogen from this substance. If two were 
removed the divalent radical CH 2 would be produced. This 
radical is known as methylene and forms a long series of com- 
pounds similar to those of methyl, but differing of course because 
of the fact that the radical is divalent. The chloride, for ex- 
ample, is CH 2 C1 2 . Under proper conditions two of these meth- 
ylene radicals will combine for the formation of the substance 
C 2 H 4 , which is known as ethylene. This is the first member of 
a long series of hydrocarbons, known as the ethylene series, 


These compounds have the property of combining directly with 
two atomic weights of hydrogen forming members of the paraffin 
series. Because of the fact that they are capable of taking up 
more hydrogen they are called " unsaturated " hydrocarbons. 
They are also able to combine directly with two atomic weights 
of a halogen per mole of hydrocarbon. Ethylene, the first mem- 
ber of the series, is formed by heating ethyl alcohol with sulfuric 
acid to a temperature of 175. The apparent reaction is 

C 2 H 5 OH = C 2 H 4 + H 2 

but it probably goes in steps, intermediate compounds involving 
the sulfuric acid being formed. It is a colorless gaseous sub- 
stance and boils at 102. Ethylene is a valuable constituent 
of ordinary coal gas, since it burns with a strongly luminous flame, 
and contributes largely to the luminosity of the gas when burned 
with the ordinary lava tip. 

Acetylene. If we imagine three combining weights of hydro- 
gen removed from methane, the trivalent radical CH would be 
produced. Two such radicals by their combination would form 
the substance C2H2, or acetylene. This compound is still more 
highly " unsaturated" than ethylene. It is the first member of 
another series of unsaturated hydrocarbons, known as the acetylene 
series. Acetylene is a colorless gas, which burns with an intensely 
luminous flame. When burned with oxygen in a properly con- 
structed burner it will produce a very high temperature. Use 
is made of this in the metal-working trades for the autogenous 
welding of metals and also for the rapid cutting of steel, etc., by 
burning a path through it. 

Acetylene is prepared by the action of water upon calcium 
carbide, CaC2, the equation being as follows: 

CaC 2 + 2H 2 = Ca(OH) 2 + C 2 H 2 

When mixed with air, in any proportions between 2.53 per cent, 
and 73 per cent, acetylene will explode. These limits are wider 
than for any other gas. A mixture of natural gas and air, for 
example, explodes only when it contains from 5 to 15 per cent, of 
methane, CH 4 . Liquid or highly compressed acetylene is explosive 
even when unmixed with air, because it is formed from its ele- 
ments with the absorption of much heat; so if decomposition 



once starts, the heat evolved will raise the temperature and make 
the reaction go faster. This is the case with all explosives. 

Acetylene is very soluble in an organic substance called ace- 
tone, (CHa^CO, which belongs 
to a group of compounds called 
the ketones, which we have not 
discussed. This solution in 
acetone is not explosive and 
cylinders containing it under 
pressure are used to supply 
acetylene for the lights on 
automobiles and for the illumi- 
nation of isolated residences. 

Fuel Gases. Gaseous fuels 
are of great importance for 
domestic and technical pur- 
poses because they are clean 
and very easily controlled. 
They are also rinding a rapidly 
increasing use as a source of 
energy in gas engines, so that 
the subject of fuel gases is one 
of no little practical impor- 
tance. Many different varieties 
of gas are in use such as natural 
gas, coal gas, wood gas, oil gas, 
and the various types of water 
and producer gases. Natural 
gas has already been discussed. 
Wood and oil gases are of rather 
limited application, so that we 
will devote our attention at 
this point chiefly to coal, water 
and producer gas. 

Coal Gas. Coal gas is made 
by heating bituminous coal in 
closed fire clay retorts, Fig. 43. It consists principally of hydro- 
gen, methane, unsaturated hydrocarbons and carbon monoxide, 
and their relative volume is in the order given. In addition to 



these gases, a large number of valuable by-products distil over; 
tar, ammonia, benzene, carbolic acid, etc., and are removed in 
the purification of the gas. Coke is left in the retorts. 

Water Gas. When steam is passed over heated coal or coke, 
carbon monoxide and hydrogen are formed, if the temperature 
is high, 

H 2 + C = CO + H 2 

or carbon dioxide and hydrogen if the temperature is lower, 
2H 2 O + C = C0 2 + H 2 

The gaseous mixture obtained 
in'- this way usually runs about 
50 per cent, hydrogen and 40 
per cent, carbon monoxide. It 
is very valuable for fuel and 
power purposes. 

The water gas process has a 
great advantage over that for 
coal gas inasmuch as in the 
former nearly all of the carbon 
is converted into gas, while in 
the latter most of it remains as 

Producer Gas. Producer gas 
is perhaps the most important 
fuel gas for industrial and power 
purposes. It is made by two 
slightly different methods. In 
the first, air is blown up through 
a deep bed of highly heated coal 
in a great stove-like furnace 
called a producer (Fig. 44). 
Where the air first enters, the 

FIG. 44. 

coal burns to carbon dioxide, 

C + 2 = CO 2 

As the latter is carried up through the bed of carbon, it is reduced 
to carbon monoxide, 

C0 2 + C = 2CO 


This carbon monoxide mixed with the nitrogen of the air and 
some unreduced carbon dioxide constitutes the producer gas 
as made by this method. 

The second method is just the same as the first except that some 
steam is blown in with the air and the reactions are those just 
given together with the ones for water gas. The reactions for 
the first method evolve heat, while those for water gas absorb it. 
By combining the two, a greater amount of the energy of the 
coal is obtained in the gas than by either alone. The gas made 
in this way contains carbon monoxide, hydrogen, carbon dioxide, 
and nitrogen. 

A gas engine using producer gas will often deliver one horse- 
power for an hour, on a pound to a pound and a half of coal used 
in the producer. The fuel efficiency of such a plant is much 
greater than that of a steam plant. 

Blau Gas. Blau gas is made by cracking petroleum by heat- 
ing it to a high temperature then cooling and compressing the- 
products to 100 atmospheres. This liquefies many of the gase 
ous substances produced. The liquid is sold in strong steel 
cylinders, which may be connected through reducing valves to 
the pipes in the building where it is to be used. The blau gas 
burns with a very bright and hot flame. 


Flames are produced by the rapid interaction of two or more 
gases whereby sufficient heat is generated to render the gases and 
their products luminous. In most of the cases to which the term 
is applied, one gas passes in a stream into a larger body of the 
other and the chemical action takes place at the point of contact 
of the two. 

In the flames with which we are most familiar, the oxygen of 
the air is one of the active gases and the other is hydrogen, 
natural gas, coal gas, or some similar mixture; and since the flame 
is located at the opening whence the gas issues into the air, we 
very naturally speak of the gas as burning, but the air has just 
as much to do with the phenomenon as the other gas. As a 
matter of fact, it is a very easy matter to burn a stream of air 
in an atmosphere of coal or of natural gas. All that is necessary 



to do to show this is to place a wide-mouthed bottle, mouth 
downward (Fig. 45), upon a suitable support, and pass into it a 
moderate stream of gas until the air is displaced, and then light 
the gas as it issues from the mouth of the bottle. Now pass a 
glass tube from which a steady stream of air is escaping up into 
the bottle. As it goes up through the flame at the mouth of the 
bottle, the air will take fire, and con- 
tinue to burn in the gas as long as the 
supply of the two is kept up. 

All flames give off light, but they 
differ enough among themselves so that 
we may conveniently divide them into 
two classes; non-luminous and luminous. 
This classification is arbitrary, because 
there is every gradation between the 
typically non-luminous flame of hydro- 
gen and that of acetylene which is the 
most brilliant of all; that of methane 
stands near the dividing line. 

When infusible non-volatile solids are 
placed in non-luminous flames, they are 
heated to incandescence and give out 
light; a mixture of 99 per cent, thorium 
oxide, with 1 per cent, cerium oxide, is 

especially brilliant under these conditions, and it is this mixture 
which is used in the Welsbach mantles for incandescent gas 

The luminous flames are nearly all produced by the burning of 
compounds of hydrogen and carbon, and there is much evidence 
in favor of the view that solid particles of carbon are formed in 
their interior, and that the greater part of the light is due to the 
incandescence of these particles. If a cold surface be placed in 
such a flame, these particles will deposit on it forming a layer 
of soot or lampblack. In fact, this is the way in which lamp- 
black is made. 

If the flame of an ordinary hydrocarbon gas, such as coal gas 
issuing from a circular opening, be examined, it will be found to 
consist of a number of conical sheaths or zones fitting one inside 
the other as indicated in the diagram (Fig. 46). In the relatively 

FIG. 45. 




non-luminous sheath "a" water and carbon dioxide are formed, 
and the heat from this radiating inward decomposes the hydro- 
carbon toward the base of the flame forming hydrogen and 
acetylene, or hydrogen, methane and acetylene; as 
the gases pass upward the temperature rises and 
finally reaches a point where the acetylene rapidly 
decomposes into hydrogen and carbon. Since this 
is a reaction which takes place with the evolution of 
heat, the temperature of the particles of carbon is 
raised even above that of the rest of the flame, and 
this is the cause of the greater part of the light. 
These changes take place in the zone "b." As the 
carbon passes up through the flame, it gradually 
burns, first to the monoxide, then to the dioxide. In 
the interior of the flame, where the supply of oxygen is limited, 
the carbon is practically all burned to carbon monoxide before 
the hydrogen burns. The luminosity of a. flame may be increased 
by putting both gases under greater pressure, or by heating the 
gas before it leaves the burner. On the 
other hand, anything which will lower the 
temperature of the flame will decrease the 
light. This may be done by holding a cold 
object in the flame or by mixing some indif- 
ferent gas such as nitrogen with the com- 
bustible gas. This both dilutes the gas and 
makes more material to be heated, and hence 
makes the flame colder. 

Bunsen Burner. Robert Bunsen, a Ger- 
man chemist, invented the burner which 
bears his name, in order to burn coal gas 
with the clean smokeless flame, which is so 
desirable for chemical purposes. A diagram 
of the burner and its flame is shown in 
Fig. 47. The gas enters the burner through 
the jet "a," and, having a high velocity, it 
draws a certain quantity of air in at the 
holes "b" in the base of the tube "f." In passing up through 
this tube, the gas and the air mix. Three well-marked sheaths 
are formed in the flame. " C " is non-luminous, and so cold that a 

FIG. 47. 


match head may be held in it for some time without igniting. 
In this, practically no chemical change takes place. "D" is 
bluish-green and very hot, especially at the upper tip, but not 
very luminous. In it the hydrocarbons are transformed into 
carbon monoxide and hydrogen. In the outer layer "e" the 
hydrogen and carbon monoxide burn to water and carbon 
dioxide. The non-luminous flame of the Bunsen burner is 
smaller and hotter than a luminous flame burning the gas at the 
same rate. 


General. Silicon, the second member of the carbon group, 
occurs very abundantly in nature in the form of its oxide and 
various silicates. This element, next to oxygen, is present in the 
earth in the greatest quantity, forming as it does about one- 
fourth of the earth's crust. 

Preparation. Silicon exists in at least two allotropic modifica- 
tions, one crystalline, the other amorphous. Some evidence 
also points toward the occurrence of two distinct crystalline 
forms, corresponding to graphite and diamond. 

Amorphous silicon may be* prepared by heating finely powdered 
quartz, which is silicon dioxide, SiC>2, with metallic aluminum 
or magnesium; the reaction which takes place with magnesium 
is represented by the following equation, 

Si0 2 + 2Mg = Si + 2MgO 

The magnesium oxide is removed by treating the resulting 
mixture with hydrochloric acid. 

Crystalline silicon is now being prepared on a large scale at 
Niagara Falls by reducing the dioxide with carbon in an electric 
furnace. It is used in the manufacture of steel. 

Si0 2 + 2C = Si + 2CO 

An alloy of silicon and iron known as ferrosilicon is also used 
in making steel. It is made by heating together sand, iron oxide 
and carbon in an electric furnace. 

Properties. Amorphous silicon is a dark brown, friable 
substance without luster, is a non-conductor of electricity, and 
has a density of 1.8. If the temperature is raised somewhat it 
takes fire and burns. The crystalline silicon is very much more 
resistant chemically and is not attacked by oxygen even at red 
heat. It is dark steel gray in color and forms opaque octahedra 
possessing a metallic luster, its density is 2.35. 



Both varieties melt at temperatures in the neighborhood of 
1,500 to a steel blue liquid which on solidifying always produces 
the crystalline form. Silicon dissolves in hydroxides of the alkali 
metals with the evolution of hydrogen and the formation of the 
corresponding alkali silicates. 

2NaOH + Si + H 2 O = Na 2 SiO 3 + 2H 2 

This reaction is used to make hydrogen for filling balloons and 
air ships. Silicon is not soluble in any single acid but is in a 
mixture of nitric and hydrofluoric acids, forming silicon tetra- 
fluoride, and is slowly oxidized by aqua regia to silicic acid. 

Silicon Dioxide. Silicon forms with oxygen one well-defined 
compound, silicon dioxide, Si02, which is analogous to carbon 
dioxide. The monoxide, the analogue of carbon monoxide, has 
been stated to exist and has been described as a brown powder, 
but as yet this has not been sufficiently confirmed. Silicon 
dioxide or silica, as it is commonly called, occurs very abundantly 
on the earth's surface both in the crystalline and the amorphous 
modifications. The purest form is rock crystal, which is colorless 
quartz. This compound when colored by different metallic 
oxides gives rise to smoky quartz and other semi-precious 
stones. Flint occurs widely distributed in nature, and is rather 
impure amorphous silica; other forms of this variety are certain 
kinds of petrified wood, opal and jasper. 

On account of its extreme hardness and its transparency, rock 
crystal finds extensive use in the manufacture of fine lenses. 
Many laboratory utensils are now made from quartz. When 
heated before the oxy-hydrogen blowpipe it behaves like glass, 
in that it becomes pasty and does not pass directly from the 
solid to the liquid state. In this condition it may be worked 
and blown into the desired shapes. These vessels are very 
resistant chemically to nearly all reagents except the alkalies, 
and consequently they may replace the more expensive platinum 
in many operations. The temperature coefficient of this fused 
silica ware is very small, and a crucible heated to redness may be 
plunged into cold water without being broken. 

Silicic Acid and Silicates. Silicates are known which corre- 
spond to an even greater variety of silicic acids than there are 


phosphoric acids. We have for example, salts of orthosilicic acid, 
H 4 SiO 4 ; meta silicic acid, H 2 SiO 3 ; diortho silicic acid, H 6 Si 2 7 ; 
dimeta silicic acid, H 2 Si 2 O 5 ; and trisilicic acid,-H 4 Si 3 O 8 . When 
a soluble silicate is treated with an acid, the solution usually 
remains clear for a time; and then there is precipitated difficultly 
soluble silicic acid whose composition is rather variable but which 
tends to approach that of meta silicic acid, H 2 Si0 3 . When 
heated, this yields the dioxide and water. 

By proper manipulation, a colloidal solution of silicic acid may 
be prepared. A colloidal solution is one which appears perfectly 
homogeneous to the eye, and which will pass through a filter, 
but which has the same boiling- or freezing-point as pure water. 
When a beam of very strong light is sent through the solution 
and a microscope focused upon its path, a multitude of particles 
are seen, so extremely minute that they cannot be perceived in 
any other way. A true solution is homogeneous even when 
examined in this way. This shows that the solute in a colloidal 
solution is not really in solution, but that it is merely very finely 
divided and in a state of suspension. 

A few of the naturally occurring silicates will be mentioned at 
this point merely to show how important to us these compounds 
are. Clay, mica, asbestos, serpentine, and feldspar, are among 
the principal silicates; while granite (a mixture of quartz, feldspar, 
and mica), basalt, prophyry, lava, pumice stone are also common 
silicate rocks. Sandstone consists principally of particles of sili- 
con dioxide cemented together with calcium carbonate, iron oxide, 
or amorphous silica from silicic acid. 

When exposed to the weather, the silicate rocks are gradually 
disintegrated under the combined action of water, carbon dioxide, 
changes of temperature, freezing and thawing, with the formation 
of clay (aluminum silicate), silica, and the carbonates of the 
metallic elements of the silicates. Under proper conditions, these 
products of disintegration will solidify into what are known as the 
sedimentary or secondary rocks, while the original silicates are 
called the igneous or primary rocks. 

Silicon Hydride. If in the preparation of silicon by means of 
metallic magnesium 1 * an excess of the metal is used, a compound 
magnesium silicide, SiMg 2 , is formed. This compound reacts 
with hydrochloric acid yielding magnesium chloride and a gas 



which on contact with the air takes fire spontaneously forming 
smoke rings of silicon dioxide. 

SiMg 2 + 4HC1 = 2MgCl 2 + SiH 4 . 

This gas is silicon hydride, SiH 4 , the analogue of methane, and 
like phosphine, if carefully purified, loses the property of spon- 
taneous inflammability, which property is probably, in this case 
also, due to the presence of some higher hydride. 

Halogen Compounds. If amorphous silicon be treated with 
chlorine, or if a current of chlorine be passed over a mixture of 
finely powdered quartz and carbon which is strongly heated, the 
compound silicon tetrachloride, SiCl 4 , is formed. Although in 
this latter reaction neither the carbon nor the chlorine alone is 
able to effect the decomposition of the silicon dioxide, their 
combined action can bring it about. 

Si0 2 + 2C + 2C1 2 = SiCl 4 + 2CO 

Silicon chloride is a colorless liquid whose boiling-point is 59. 
It fumes strongly in the air since water readily decomposes it 
into silicic acid and hydrogen chloride. 

If instead of chlorine, hydrogen chloride is passed over amor- 
phous silicon, a compound of the composition SiHCl 3 is formed 
and by analogy is called silico-chloroform. 

Silicon unites with the other halogens to form similar com- 
pounds. The most important of these is the fluoride which is a 
gas at ordinary temperatures and is readily formed by treating 
silica or a silicate with hydrofluoric acid, 

4HF + Si0 2 = SiF 4 + 2H 2 O 

Silicon fluoride is decomposed by water, hence some dehydrating 
agent such as sulfuric acid must be added to remove that formed 
in the reaction. Silicon fluoride is most conveniently prepared 
by mixing together silicon dioxide and calcium fluoride, and 
adding concentrated sulfuric acid to the mixture. The reaction 
with water is unlike that of the other halogen compounds. 
Instead of silicic acid and hydrogen fluoride, we have the 
complex hydrofluosilicic acid formed according to the equation, 

3SiF 4 + 4H 2 O = 2H 2 SiF 6 + Si(OH) 4 



The lead salt of this acid is used in the electrolytic purification of 
lead (see p. 417). 

Carborundum. Carborundum is the trade name for silicon 
carbide, SiC. It is chemically resistant, so hard it will scratch 
ruby, and is used very largely as an abrasive. It is not attacked 
by acids but is gradually decomposed by fused alkalies forming 
the respective carbonate and silicate. 

Carborundum is prepared on a large scale at Niagara Falls 
by heating together in an electric furnace (Fig. 48), carbon, 
sand and sawdust with common salt as a flux. The electric cur- 
rent is passed through the carbon core, T (Fig. 48), until the mix- 

v T 

After Run. 

B = Sand and Carbon. 

D Crystalline Carborundum. 

E - Graphite and > 

F - Carbon Core 

= Amorphous Carborundum. 

r?- Copper Electrodes. 

5= Sand and Carbon 

T** Granular 

V = Carbon Powder. 

ture of carbon, sand, etc., has been heated sufficiently. The 
furnace is then allowed to cool when a cylinder of crystallized 
carborundum will be found surrounding the carbon core. It 
is then crushed, treated with sulfuric acid to remove impurities 
such as iron oxide, washed and graded. The product is then 
worked up into abrasive wheels, hones and the like. For this 
purpose the carborundum is mixed with a suitable binding 
material such as china clay and feldspar, molded and baked at a 
high temperature. 


General. Boron is very similar to silicon in properties, both 
of the free element and of its compounds. The formulas of these 
latter, however, together with some other considerations connect 
it quite closely with certain of the true metals. It occurs in 
nature as boric acid, H 3 BO 3 , the salts of this acid, and as the 
anhydride or boron trioxide, B 2 O 3 . From these it will be seen 
that unlike silicon, boron is trivalent. The free element may be 
prepared in both the crystalline and amorphous varieties as in 
the case of silicon, and the method is much the same as for 
that substance, viz., by the reduction of the oxide with a metal. 

Fused boron has recently been prepared by reducing the tri- 
oxide with magnesium, and then heating the product to 2000 
2500 in an electric furnace when it melts. 

B 2 3 + 3Mg = 2B + 3MgO 

The product is black, amorphous and very hard, scratches 
everything except diamond, but is rather brittle. When cold, 
it is a very poor conductor of electricity, but at 400 it conducts 
2,000,000 times as well as at room temperature. Because of its 
great temperature coefficient, it is well adapted to the making 
of devices for the measurement of temperature and for the 
regulation of electrical machinery. 

Boric Acid. Boric acid, or boracic acid as it is commonly 
called, is formed by the action of water on the anhydride, boron 
trioxide. The acid corresponding by analogy to orthophosphoric 
acid, is H 3 BO 3 , and is called orthoboric acid. While this acid 
exists free, nearly all the known salts are formed from acids 
derived from H 3 B0 3 by the loss of water. In nature it occurs as 
somewhat yellowish scales. It is very soluble in hot water but 
much less so in cold. The aqueous solution reacts faintly acid. 

On being heated, the acid passes first into water and the tri- 
oxide, which at a higher temperature melts to a glass-like mass 


BORON 261 

which is capable of dissolving metallic oxides. Upon this 
property depends the use of boric acid as a flux in hard soldering. 

While the anhydride is fairly resistant to heat, the acid itself 
is volatile with steam. One of the methods for obtaining the 
acid depends upon this fact. In the volcanic districts, vapors 
containing boric acid escape from the earth. These are passed 
into water, condensed, and the boric acid finally becomes concen- 
trated enough to crystallize out at lower temperatures. When 
boric acid is introduced into a flame it volatilizes and imparts a 
green color to the flame. 

Of the many boric acids, we shall mention only the metaboric, 
formed from H 3 BO 3 by the loss of one mole of water, according 
to the equation, H 3 BO 3 = HB02 + H 2 O; and the tetraboric 
formed according to the equation 4H 3 BO 3 = H 2 B 4 O7 + 5H 2 O. 
In any water solution of boric acid, the ions of these different 
acids exist in equilibrium. 

The sodium salt of tetraboric acid, Na 2 B 4 07-10H 2 O or borax, 
occurs in California and in others of the western states and is 
also manufactured by decomposing the naturally occurring 
calcium borate, Ca 2 B 6 On-5H 2 0, with sodium carbonate and 
bicarbonate and recrystallizing. Borax contains either 5 or 10 
moles of water of crystallization depending upon the temperature 
at which it is crystallized. It effloresces in the air and, if heated, 
swells up, then melts to a glassy substance which dissolves metallic 
oxides similarly to boric oxide. Some of these solutions have 
distinctive colors, and are used in blowpipe analysis and in 
making enamels. Because of the weakness of boric acid a solu- 
tion of borax is hydrolized enough to be alkaline. On this 
account it is used to a slight extent in households to soften water, 
but this is a very costly method. Borax and boric acid have 
fairly strong antiseptic properties and are used in medicine and 
as preservatives. 

Other Compounds. The chlorides and other halogen com- 
pounds of boron may be prepared in the same way as the corre- 
sponding compounds of silicon and are similar in their properties. 
Since boron is trivalent their general formula is BA 3 , where 
A represents the halogen anion. The fluoride when dissolved in 
water forms the compound, HBF 4 , which is very much like 
hydrofluosilicic acid in its properties. 


Boron readily combines with nitrogen to form boron nitride, 
BN, so unless special precaution is taken to exclude the air during 
the preparation of the free element, the latter is likely to be con- 
taminated with the nitride. This compound is a white powder 
which when heated with water reacts to form ammonia and boric 

BN + 3H 2 O = B(OH) 3 + NH 3 


The gaseous substances helium, neon, argon, krypton, xenon 
and niton form an interesting group of non-metallic elements 
which is chiefly remarkable for the fact that no compound of 
a member of this group has ever been prepared. Another in- 
teresting point in connection with these substances is that the 
element helium is certainly one of the products of the sponta- 
neous transmutations which the radioactive element radium 

The molecular weights of these elements may be easily ob- 
tained from their gaseous densities but since they will not enter 
into combination with any substances, the question as to the 
number of atoms per molecule and hence their atomic weight is 
not easily answered. But in spite of their inactivity we can 
arrive at a reasonable conclusion through the following line of 
argument. ' 

The quantity of heat necessary to raise the temperature of a 
gram mole of a substance 1 is called its molecular heat. If the 
volume is kept constant during the heating, it is known as 
the molecular heat at constant volume. Now it is an easy 
matter to determine both the molecular and atomic weights of 
mercury, and these turn out to be 200 in each case. In addition, 
the molecular heat at constant volume of mercury vapor is 13 
joules. The molecular heat at constant volume of gases having 
two atoms per molecule such as oxygen, nitrogen, hydrogen, 
carbon monoxide, etc., lies between 20 and 21 joules, while that 
for substances having three atoms per molecule is 28 to 33 joules. 
The molecular heat at constant volume for the members of the 
argon group is 12 to 13 joules. We may therefore safely con- 
clude that there is one atom per molecule of these elements just 
as is the case with mercury, and hence that their atomic and 
molecular weights are identical. 



Helium. The element helium has an atomic and molecular 
weight of 4.00. It is found in very small quantites in the air. 
It is also found in the gases of many mineral springs and is pres- 
ent in relatively large quantities in natural gas. The gas at 
Dexter, Kansas, contains as high as 1.84 per cent, of helium. 
Outside of natural gas, the principal source of helium is the rare 
minerals containing the radioactive elements radium, uranium, 
thorium, etc., from which it may be obtained by simply heating 
the mineral in a vacuum or by dissolving it in dilute sulphuric 
acid or fused potassium acid sulfate. The helium is simply 
enclosed in the mineral and is not in combination. From the 
fact that it occurs in those minerals which contain radium, etc., 
it is probable that it is formed in the mineral by the transforma- 
tion of the radioactive elements. The same source is indicated 
for the helium of natural gas by the fact that it contains niton 
(see p. 265). 

Helium was the last gas to be liquefied. This was done in 
July, 1908, by Kammerlingh Onnes. It is a clear, colorless liquid 
boiling at 268. 5C. or 4.5 above absolute zero. Since its 
critical temperature is 268. 0C. or + 5 A. it is a very difficult 
substance to liquefy. When the pressure under which it boiled 
was reduced to 1 cm. its boiling-point was lowered to 3 above 
absolute zero. This is probably about the lowest temperature 
which we will be able to reach unless some more easily volatile 
liquid than helium is discovered. 

Helium under diminished pressure, glows brilliantly upon the 
passage of an electric discharge. Upon examining the light with 
a spectroscope the spectrum is found to consist of a number of 
bright lines, one in the yellow being especially strong (see Frontis- 
piece). Many years before helium was discovered on the earth 
this line was observed in the spectrum of the sun and ascribed to 
a then unknown element which was called helium from the Greek 
name for the sun "helios." 

Neon. Neon has an atomic and molecular weight of 20.2. 
It is found in the air in slightly larger quantities than helium. 
It occurs along with helium in the gases of certain mineral springs 
and in natural gas. Its boiling-point is 246. It is easily 
liquefied in a bulb surrounded by liquid hydrogen which boils 
at 253, and is separated from helium in this way. Neither 


helium nor neon is absorbed by cocoanut charcoal to any great 
extent, and a mixture of the two gases may be easily obtained 
from natural gas by absorbing the other gases in cocoanut char- 
coal cooled with liquid air. Recent experiments indicate that 
neon may be a mixture of two gases one having an atomic weight 
of 20 and the other of 22. 

Argon. The atomic and molecular weights of argon, are 
39.9. It occurs in nature in rather large quantites being pres- 
ent in the air to the extent of about 1.3 per cent, by weight or 
0.937 per cent, by volume. It boils at 186 and freezes at 
189. Argon was the first member of this group to be dis- 
covered, being found in 1894 in atmospheric nitrogen by Lord 
Rayleigh and Professor Ramsay who were seeking an explana- 
tion for the fact that atmospheric nitrogen was more dense than 
that from chemical sources. It is separated from nitrogen by 
combining the latter with metallic lithium, calcium, or mag- 
nesium at a high temperature, nitrides being formed. It is now 
being used in the place of nitrogen in some of the gas-filled tung- 
sten lamps. Its spectrum is shown in the Frontispiece. 

Krypton and Xenon. These gases occur in the air in very small 
amounts. They are obtained mixed with argon, helium, and 
neon by the removal of the nitrogen, etc., as indicated above. 
They are separated in the pure state by the fractional distilla- 
tion of the crude argon. Krypton has an atomic and molecular 
weight of 82.92 and boils at 152. Xenon has a molecular and 
atomic weight of 130.2, boils at - 109. 

Niton or Radium Emanation. Radium undergoes a slow 
spontaneous transmutation and very gradually gives off an inert 
gas of the argon group which is called niton or radium emanation. 
Although this is .chemically inactive, it is highly radioactive, 
(see p. 495) and changes into helium and other products at such 
a rate that one-half of it disappears every 3.8 days. Since it is 
formed so slowly from radium and decomposes so quickly, there 
is never very much of it available for experimentation at any one 
time, and Sir William Ramsay who has studied it most fully 
has displayed great skill in its investigation. It has been shown 
in spite of the fact that it is formed from one element and changes 
into another, that it is fully entitled to be considered a distinct 
element. It has its own characteristic spectrum and set of other 



properties. Its molecular and atomic weights have been deter- 
mined by Ramsay by weighing less than 0.0001 c.c. of the gas 
on a balance sensitive to less than H 5 0,000 rngrm. These 
weights are doubtless identical and are 222.4. It boils at 
62 and freezes at 71. Niton is so highly radioactive that 
it shines in the dark, and gives out, in 
comparison with its mass, enormous 
quantities of heat. When brought in 
contact with chemical compounds, it 
very frequently decomposes them, but 
will not itself combine with anything. 
Niton is present in extremely minute 
quantities in air, soil, rocks, ground 
water, and the gases that issue from the 
earth. The quantity present in rocks is 
proportional to their radium content 
and furnishes the easiest method of de- 
termining the amount of radium in a 
rock. The rock is dissolved, the niton 
boiled off and passed into a charged 

gold leaf electroscope (Fig. 49) ; the quantity of niton and there- 
fore of radium is estimated from the rate at which the gold leaf 
of the electroscope collapses. 

The following table giving the per cent, by volume of the 
inert gases in the air is based upon an estimate by Sir William 


Helium 0.00040% 

Neon 0.00123 

Argon 0.9370 

Krypton 0.0094 

Xenon 0.0011 

Niton.. . Trace. 

FIG. 49. 


General. In the classification of the elements it has been 
customary for a long time to divide them into the metals and 
the non-metals. We have been considering the latter in the 
earlier part of this book, and it now remains to discuss the 
chemistry of the metals. The division of the elements into these 
two groups is very convenient, but is entirely arbitrary as is 
shown by the fact that certain elements have properties in 
common with each group and cannot be definitively located in 

Characteristic Physical Properties of Metals. All metals 
when not powdered reflect so much of the light which falls upon 
them that they have an appearance called a metallic luster. 
Most are silvery white, but a few, such as gold and copper, absorb 
enough light to show distinct colors. But everything with a 
metallic luster is not a metal. The non-metal tellurium shows 
it, as do many compounds, especially the sulfides. 

All metals are much better conductors of electricity than 
electrolytes are, and, unlike the latter, are not decomposed by the 
passage of the current, hence they are called conductors of the 
first class. But several of the non-metals, carbon, boron, selenium, 
tellurium and many sulfides are also conductors of the first 

The metals are better conductors of heat than the non-metals, 
but in this respect there is not as much difference between the 
two groups as in the case of the conduction of electricity. 

Most, but not all of the metals, are malleable and may be rolled 
or beaten into thin plates. Gold excels all the others in this 
respect, and may be beaten into leaves not more than 0.0001 
m.m. in thickness. 

Most metals are very tenacious and will resist a considerable 
force tending to pull a given piece apart. This property if 



associated with the malleability mentioned above renders a 
metal ductile, so that it may be drawn into wire. Steel is the 
most tenacious of the metals, and gold the most ductile. 

When metals, dissolve in most solvents they lose all of their 
metallic characteristics, and in general form salts, but they may 
dissolve in one another when fused together and retain the 
characteristic properties of metals. The solutions of metals 
in one another are called alloys. As the term alloy is now used, 
one or more of the constituents may be a non-metal, but the whole 
must present the general metallic properties. Alloys which 
contain mercury are called amalgams. Some of these are very 

Characteristic Chemical Properties of Metals. The most 
noticeable chemical property of the metals is that they form 
simple cations, but this property is not confined exclusively to 
them; because the non-metals, hydrogen and tellurium also form 
these ions. Closely connected with the tendency of the metals 
to form cations is the fact that their hydroxides are generally 
bases while those of the non-metals are almost invariably acids; 
but some of the hydroxides of the metals are very weak bases, and 
a few act both as bases and acids. From the above facts, it may 
be seen that there is no sharp distinction to be drawn between 
metals and non-metals, so the division is purely arbitrary; but 
is nevertheless convenient and practicable because in the great 
majority of cases no hesitation is felt in making the classification. 

Distinction between Metals and Non-metals. If an element, 
when in the solid or liquid state, does not have a metallic luster 
and is not a conductor of electricity, it certainly is a non-metal. 
If it has the luster and conductivity, it may be either a metal 
or a non-metal and we must turn to its chemical properties for 
help in its classification. If it forms simple cations, it is a nietal; 
but if it gives no simple cations and is nearly always found in 
anions, it is a non-metal. After all aids to classification have been 
applied, a few elements on the dividing will still be in dispute as 
to whether they are metals or non-metals, but the great majority 
will be definitely located. 

Classification of the Metals. A study of the properties of the 
metals soon shows that they may be divided into groups of more 
or less closely related elements much as was done with certain of 


the non-metals. The relationship between these elements is 
rather tangled, however, and some of them have properties which 
would cause them to be placed in each of several groups and any 
arrangement which we make of them is more or less arbitrary. 
There is, however, a natural system for the classification of the 
elements which is called the periodic system and which includes 
both the metals and the non-metals. This system is based upon 
the observation that if the elements be arranged in the order of 
their atomic weights, similar elements recur at regular inter- 
vals; and if the series be broken up into periods such that each 
period begins with a member of a definite family, the second 
position in each period will be filled by a member of a closely 
related group of elements. The same is true for the third and 
following positions in the series. This may be illustrated as fol- 
lows; arrange the elements in the order of their atomic weights 
and we have, 

Position, 1, 2, 3, 4, 5, 6, 7, 8 

Period No. 1. He. Li. Be. B. C. N. O. F. 

Period No. 2. Ne. Na. Mg. Al. Si. P. S. Cl. 

Period No. 3. A. K. Ca. Sc. Ti. V. Cr. Mn. 

Choosing the members of the helium family as the elements 
which shall begin our periods, we find that the elements listed 
here fall into three periods of eight members each, and that the 
corresponding positions in each period are occupied by analogous 
elements. We can appreciate this especially for the last four 
members of the first two periods, since we are familiar enough 
with the properties of carbon and silicon to see their close rela- 
tionship and similarly for nitrogen and phosphorus, oxygen and 
sulfur, fluorine and chlorine. That this same sort of relation- 
ship is 'continued over into the other periods and holds for the 
remaining positions in these periods, will be evident as we 
become more familiar with the properties of the elements. For 
example, lithium, sodium, potassium, etc., which occupy the 
second positions in the periods, form a part of a group of similar 
and very closely related metals as we shall soon see. In fact 
the relationship is fully as close as that between the halogens. 

If the elements arranged in series in the order of their atomic 
weights be broken up into periods in the way given above and 
these periods be placed vertically under one another, we ob- 


tain a table of the periodic classification or system of the ele- 
ments. Such a table is given below. It is somewhat of a 
modification of one prepared by the Russian chemist, Mendele- 
jeff in 1869, and very similar to a table published in 1870 by 
Lothar Meyer, a German chemist, who developed the idea 
independently of Mendelejeff. 

An inspection of this table will show that chemically similar 
elements are found in vertical columns except in the case of the 
very last group in which the similar elements are found side by 

In preparing this table, the principle of arrangement in the 
order of atomic weights has been violated in two cases. Argon 
with an atomic weight of 39.9 has been placed before potassium 
with an atomic weight of 39.1 and tellurium is put before iodine 
although it has the larger atomic weight. The reason for the 
first change is that potassium does not belong in the group of inert 
gases, but in with the other alkali metals in which group the 
argon would be entirely out of place while it goes very well 
with the other gases of the helium group. Tellurium and iodine 
are interchanged because the tellurium evidently belongs to the 
sulfur group and iodine to the halogen. Many researches have 
been undertaken to see if there was not some error in the deter- 
mination of the atomic weight of either tellurium or iodine but 
they have led to the conclusion that tellurium actually has a 
higher atomic weight than iodine. 

Another defect in the table is that there is no good place for 
hydrogen. Some have placed it at the head of the halogen 
family and others as the first member of the sodium group. But 
neither these nor any other suggestions which have been made 
seem to fit the case and the question of the proper location of 
hydrogen in the periodic system remains an open one. 

There are other imperfections in the table which will become 
evident as we go along, but on the whole it furnishes a natural 
and very useful system for the classification of the elements, and 
it seems not unlikely that its apparent irregularities and defects 
will finally be found to have a great significance and to be very 

Although each of the first eight vertical columns in the table is 
made up of very similar elements, a knowledge of their properties 



s 2 

Oi T3 *^ 


O * 





^ 1 s 

t 2 2 

2 - T. 


00 ^ 

jg , 8 

<0 3 oo 

fe rt o 





CO * l> <N 




OS . O ^ I-H 








O N 
N O OS t 



CO S CD N 3J 00 





"5 -! . N 

M - ZB O . 

CD o S ^ ; P 




o >o cs ,_ 






^ M \4 co ^ g 

^& ^ g ^g 3 - Np 




CO 5 t>i CO N & 





c5 ft' N u ^* H 





^^g^^os' SO' 




- 5 " j s * " 5 i e - i 






CC ^* ^ ^ * 5O O 

2 . I s - *- *-* c* cv 
* ^ " fc^ ^ * fc ^ 



^. ^ S ^ 2 g : 




3 M U S ? 0* < \ 


O OS s N N 
(N N ' ' <N 



* e m oo J2 

W <" W " -2 



---*-.--- S 2 






soon shows that each group is composed of two sub-groups, and 
that the members of each sub-group are found in the alternate 
rows or periods. This is brought out in the table by setting the 
rows or periods alternately to the right and left of the group. 
As a rule, the members of each sub-group are more nearly like 
one another than they are like the members of the other sub- 
group within the family. The first member of each group usually 
has properties in common with both of the sub-groups and forms 
a connecting link between them. 

Group VIII is a rather exceptional group in that it is composed 
of three sets of elements, each set having three members whose 
atomic weights, specific gravities and general properties are 
very similar. In this group, the elements which are most alike 
are found side by side, although there is also a distinct tendency 
toward similarity in the vertical columns. 

It will be noticed that these sets of three elements come at the 
end of the third, fifth, and ninth series or periods of the elements. 

A blank space in the table indicates that no element is 
known whose atomic weight and properties would give it this 
position. It is very probable that some of these elements will 
be discovered in the course of time. In the meantime it is 
possible to make a fairly definite prediction as to the properties 
of an element from those of the elements surrounding it in the 
table. This was done with surprising accuracy by Mendelejeff 
for the elements of scandium, gallium, and germanium which 
were unknown at the time of the publication of the table. 

Not only are the chemical properties of the elements regularly 
recurrent or periodic functions of the atomic weights, but the 
same is also probably true of all measurable physical properties 
with the exception of the specific heat which is inversely propor- 
tional to the atomic weight for all solid elements of an 
atomic weight higher than 35. This is very clearly shown by 
Fig. 50 which represents the relation between the density of the 
elements in the liquid or solid state and the atomic weight 
and also that between the absolute temperature of fusion and 
the atomic weight, the former being shown by the heavy 
lines and the latter by the dotted lines. Most other physical 
properties yield similar curves and indicate clearly that the 
properties are periodic functions of the atomic weights. 



It would be very interesting to know the proportions in which 
the various elements are present in the earth as a whole, but this 
is manifestly impossible since we have no way of getting at the 
composition of its interior. We can, however, get the analysis 
of the air, the ocean, and the rocks composing the outer 10 miles 
of the earth's crust, the lithosphere as it is called, and the follow- 
ing table (p. 274) by Clarke of the United States Geological 
Survey gives the results of some thousands of analyses. 

FIG. 50 

All the abundant elements have an atomic weight below 56 
and the heavy metals appear to be present in very small quan- 
tities. The mean density of the earth is about twice that of the 
lithosphere and hence it has been supposed that the heavier ele- 
ments have collected in the interior. Others ascribe the high 
density to the compression due to the weight of the overlying 
strata. This much is certain, however, that {he elements cannot 
all be present in equal quantities by weight since their average 
density is greater than the mean density of the earth. 





(93 per cent.) 

(7 per cent.) 

Average in- 
cluding nitro- 


47 33 

85 79 

50 02 

Silicon . 

27 74 

25 80 


7 85 

7 30 



4 18 

Calcium. . . 



3 22 


2 24 


2 08 


2 46 

1 14 

2 36 




2 28 



10 67 





Ca'rbon. . . 






2 07 

























All other elements 







Group I of the periodic system is composed of the mono- 
valent alkali metals, lithium, sodium, potassium, rubidium and 
cesium, and a distinct sub-group consisting of copper, silver, and 
gold. The alkali metals resemble one another closely, and 
hence may very properly be considered together. In many 
respects their properties are typical of the metals and a discussion 
of this class of elements may profitably begin with them. Cop- 
per, silver, and gold, however, differ so markedly from the alkali 
metals and have so many properties similar to those of metals 
belonging to higher groups, that it seems wise to leave their 
discussion to a later point. 

Some of the more prominent properties of the alkali metals are 
given below in tabular form for sake of comparison. 


Atomic weight 





6 94 





23 00 


97 6 


Potassium . . . . 





Rubidium. . . ... 






132 81 




It will be noticed that, as the atomic weight increases, the melting- 
and boiling-points decrease and the density increases. Sodium, 

1 A list of the members of each group will be given at the beginning of the 
discussion of the group. To save space, this will be put on one line. The 
member of the group having the smallest atomic weight, which is also the 
connecting link between the two sub-groups, will be placed about the middle 
of the line with the members of the two sub-families at its right or left ac- 
cording as they fall into the right or left hand column in the group in the 
periodic system, 



however, is somewhat irregular in its density, a fact which may 
be connected with the position of sodium in the periodic system 
as the first member of the copper sub-group since these metals 
all have high densities. This individuality of sodium seems 
to extend to its chemical properties since potassium, rubidium, 
and cesium resemble one another much more closely than they 
do sodium. 

The alkali metals are chemically the most active of this class 
of elements and increase in activity as their atomic weight 
increases. They will all decompose water, lithium rapidly 
and cesium with the greatest violence. In each case, hydrogen 
and the hydroxide of the metal is formed. These hydroxides 
are soluble and are strong bases, and it is this fact that gives to 
the sub-group the name, the alkali metals. 

None of these elements is ever found free in nature, a fact 
which might be anticipated from their great chemical activity. 
Sodium and potassium are among the more common elements 
and do not differ much in their abundance, but owing to the fact 
that enormous quantities of sodium chloride have been deposited 
in the salt beds which are found widely distributed over the 
earth's surface, while only a few notable deposits of potassium 
salts are known, the sodium compounds are much more avail- 
able and consequently much cheaper than the compounds of 
potassium. Since the sodium salts are fully as good as the potas- 
sium for all purposes which do not demand the presence of 
potassium ion, much larger quantities of the sodium compounds 
are manufactured and used than of potassium compounds. For 
this reason, the detailed discussion of the alkali metals will begin 
with that of sodium. 


Occurrence. Enormous quantities of sodium compounds, 
chiefly the chloride, are found in sea water, and in places where 
some of the sea water has been cut off from the rest of the ocean 
and evaporated, great beds of salt have been formed. So the 
supply of this important substance is practically without limit. 

Other important sodium compounds occur in nature in con- 
siderable quantities. These are the carbonate Na2COs and the 



bicarbonate, NaHC0 3 , the nitrate, NaN0 3 , cryolite, Na 3 AlF 6 , 
borax, Na 2 B 4 O 7 , and many of the silicates. 

Preparation. Metallic sodium was first prepared by Davy 
in 1807 by the electrolysis of fused sodium hydroxide, the method 
which is used at the present day. Sodium hydroxide is melted 
in an iron vessel (Fig. 51) and electrolyzed between iron or 
nickel electrodes. Sodium and hydro- 
gen are liberated at the cathode and 
oxygen at the anode; the metallic 
sodium rises to the top of the bath 
and collects under an iron bell-like 
vessel from which it is dipped from 
time to time. The sodium hydroxide 
is replaced as it is used up, thus 
making the process continuous. 
Thousands of pounds of sodium are 
now being made yearly by this process 
at Niagara Falls. 

Many attempts have been made 
to prepare sodium by the electrolysis 
of fused sodium chloride, but the ap- 
paratus used has always been so short lived that the processes 
have not been successful. 

Properties. Sodium is a soft metal with a brilliant silvery 
white luster which is almost instantly lost upon exposure to the 
air owing to the formation of a film of oxide or hydroxide. It 
melts at 97.6 and boils at 878. The vapor has a blue color 
and is exceedingly active chemically. This makes it difficult 
to obtain the molecular weight of sodium, but the results indi- 
cate that its value is 23, the same as the atomic weight. 
The metal is soluble in mercury, forming sodium amalgam. 
When this contains more than a small amount of sodium, it is 
solid. The amalgam is less active chemically than the metal 
itself, and finds extensive application, because, for many purposes, 
metallic sodium is too vigorous. 

The principal chemical property of sodium is its great tendency 
to pass over into the ionic state. In doing this, the sodium takes 
up positive electricity and hence is a strong reducing agent. 

Metallic sodium reacts vigorously with water, with the forma- 

FIG. 51. 


tion of sodium hydroxide and hydrogen. If more than a very 
small piece of sodium is used, a violent explosion is liable to 
occur. Ordinarily, the hydrogen evolved does not become suffi- 
ciently heated to take fire, but if the motion of the globule of 
sodium be diminished by placing a piece of filter paper on the 
surface of the water, the hydrogen catches fire and burns with a 
yellow flame, the color being due to the sodium. 

Moist air acts rapidly upon sodium, but dry air at ordinary 
temperatures works very slowly. Large quantities of the 
metal are preserved in soldered tin cans or closely stoppered 
bottles. Small quantities of the metal are best kept under a 
petroleum oil which is without action on the metal. 

When heated in the air sodium burns with a yellow flame 
forming the oxide and peroxide. The larger part of the sodium 
manufactured is used in the preparation of the peroxide and of 
the cyanide. 

Sodium Hydride. When sodium is heated to 340 in an 
atmosphere of hydrogen a crystalline compound, NaH, is formed. 
This dissolves in water with the formation of sodium hydroxide 
and twice as much hydrogen as an equivalent quantity of the 
metal. Because of the great amount of hydrogen liberated per 
unit weight of the substance, it has been proposed to use it as 
a source of hydrogen for filling balloons. In marked distinction 
from most of the hydrogen compounds of the non-metals, sodium 
hydride has no acid properties. 

Sodium Chloride. Sodium chloride or common salt, NaCl, 
may be prepared by the interaction of moist chlorine and metallic 
sodium or by neutralizing sodium hydroxide with hydrochloric 
acid and evaporating. 1 

However, sodium chloride is never manufactured as indicated 
above, because of the fact that it occurs in nature in large quanti- 
ties in easily worked beds of salt. From these the salt may be 
taken by ordinary mining operations when it is known as rock 
salt; or it may be obtained by boring down into the salt bed and 
running water down into the boring, allowing it to stand until it 
is saturated with salt, then forcing or pumping out the brine 
and evaporating it to secure the salt. Large quantities of salt 

1 It is remarkable that dry chlorine will not act on sodium even at the 
melting-point of the latter. 


are obtained by the evaporation of brine from salt springs or from 
the ocean. Water from the ocean is usually concentrated by 
exposure to the heat of the sun in shallow basins constructed 
near the shore. The salt so obtained is called solar salt. The 
principal impurities contained in salt beds and brines are calcium 
sulfate and chloride, magnesium chloride and the bromides. 

To remove these, the salt must be purified by crystallization. 
This cannot be done by preparing a hot saturated solution and 
allowing it to cool, since the salt is nearly as soluble in cold as in 
hot water. The salt is purified by boiling the solution which 
quickly causes the greater part of the calcium sulfate to pre- 
cipitate, and then concentrating until most of the sodium chloride 
has crystallized out; but stopping the operation before the other 
impurities have begun to separate. Magnesium or calcium 
chlorides are particularly objectionable, since they cause the 
salt to deliquesce under ordinary atmospheric conditions. 
Pure sodium chloride will not deliquesce except in the very 
dampest weather of the summer. 

The very pure sodium chloride required for many chemical 
purposes is obtained by dissolving the salt in water, and saturat- 
ing the solution with hydrogen chloride which precipitates out the 
greater part of the salt. The explanation of this will be given 

Aside from its use in food and as a preservative, sodium chloride 
is the source of practically all other sodium compounds and of 
the greater part of the chlorine and chlorine compounds of com- 
merce. For these reasons it is a very important substance and 
is used by the millions of tons yearly. 

Sodium Bromide and Iodide. Sodium bromide and iodide 
are very similar to the chloride. 

The solubility of the bromide is greater than that of the 
chloride and is in turn exceeded by that of the iodide, so it may 
be seen that the solubility increases as the atomic weight of the 
halogen increases. 

Both salts are used in medicine and the bromide in photography. 

Sodium Oxides. Sodium monoxide, Na0, may be prepared 
by heating, metallic sodium with sodium nitrate or nitrite in the 
absence of air. The equation for the latter reaction is as follows: 
2NaN0 2 + 6Na = 4Na 2 + N 2 


It is a gray substance which reacts violently with water to form 
the hydroxide NaOH. 

When heated in an excess of air to 300, sodium is oxidized to 
sodium peroxide, Na 2 O2. To carry out the process, the sodium in 
thin pieces is placed in aluminum trays which are slowly passed 
in one direction through an iron tube heated to about 300 while 
a current of dry air, purified from carbon dioxide, is passed in 
the opposite direction. This brings the air from which the 
oxygen has largely been removed in contact with the fresh sodium 
and so moderates the reaction, while the nearly oxidized sub- 
stance comes in contact with the fresh air and is fully oxidized. 

We have here one reacting substance moving in one direction 
and the other in the opposite. This arrangement is often adopted 
in chemical work and is known as "the principle of counter 
currents." The sodium peroxide so prepared finds extensive 
use as an oxidizing agent and as a source of oxygen, since it 
reacts with water w.ith the formation of sodium hydroxide and 
hydrogen peroxide which then breaks down into water and 
oxygen. Fused sodium peroxide, known as oxone, is put on the 
market for this purpose. A small quantity of a copper compound 
is added to the peroxide to catalytically decompose the hy- 
drogen peroxide. 

Sodium Hydroxide. Sodium hydroxide may be prepared by 
the action of the metal upon water, and is so prepared when a 
very pure article is required. The hydroxide obtained by this 
method is of course too expensive for most purposes, and the 
bulk. of the substance must be prepared in other ways. It may 
be made in a purely chemical way by adding calcium hydroxide 
in suspension in water, milk of lime as it is called, to an 8 to 10 
per cent, solution of sodium carbonate and boiling for a time. 
Difficultly soluble calcium carbonate is precipitated and a solu- 
tion of sodium hydroxide is left : 

Na 2 C0 3 + Ca(OH) 2 = CaC0 3 + 2NaOH 

Sodium hydroxide is also prepared by the electrolysis of a 
solution of sodium chloride using mercury for the cathode. 
Under these conditions, sodium amalgam is formed which 
is removed from the electrolytic chamber and allowed to react 
with water when hydrogen and sodium hydroxide will be formed. 



The mercury freed from the sodium is returned to the electro- 
lyzing chamber and retransformed into the amalgam. At the 
graphite anodes, chlorine is evolved. This process is carried out 
at Niagara Falls and other places where power is very cheap. 

Another successful electrolytic process is that using what is 
known as the Townsend cell. This is shown in Fig. 52. The 
anode is graphite; D and D are asbestos diaphragms painted 
with a mixture of iron oxide, asbestos fiber, 
and colloidal ferric hydroxide. The 
cathodes are two in number, and are made 
of perforated plates of iron directly in con- 
tact with the diaphragms. The anode 
chamber is filled with brine which perco- 
lates slowly through the diaphragms and 
sinks to the bottom of the cathode cham- 
bers, the remaining space in these cham- 
bers being filled with kerosene. From 
these, the brine passes off through the 
goose necks, B. When the current is 
passed through the cell, chlorine is given 
off at the anode, and hydrogen and sodium 
hydroxide appear at the cathode. Since 
hydroxyl is an anion, it tends to travel 
toward the anode, but the brine is made to 
percolate through the diaphragms rapidly 
enough to sweep back the hydroxyl and hence prevent what 
would otherwise be a serious loss. The cathode solution con- 
tains both sodium hydroxide and chloride. As this is evapo- 
rated, the chloride becomes practically insoluble, and may be 
readily separated from the hydroxide. 

Sodium hydroxide is a deliquescent white substance which 
is exceedingly soluble in water. It is a very strong base and is 
generally used when such a substance is required as in the manu- 
facture of soap, paper-pulp, and in many other chemical processes. 
It is also commonly used in the laboratory when a solution 
containing the hydroxyl ion is needed. 

Sodium Carbonate. Sodium carbonate, Na 2 C0 3 , has so many 
uses that it is almost indispensable to our civilization. It is 
found in. nature in deposits and in solution in the lakes of the 

FIG. 52. 



drier parts of the earth's surface. The United States has a 
number of such lakes, Mono Lake and Owens Lake in California, 
for example. These two lakes contain over a hundred million 
tons of the carbonate. It can be extracted cheaply but high 
transporation charges prevent its general use. 

The ashes of sea plants contain sodium carbonate and this was 
formerly the chief source of the salt. At that time this com- 
pound was at least ten times more expensive than at present and 
this raised the price of many other things, notably glass and soap. 

The oldest practical method for the preparation of sodium 
carbonate from the chloride is that known as the Le Blanc 
process. It was invented by Le Blanc in 1791. The method 
involves essentially three steps. The first consists in the con- 
version of sodium chloride into sulfate, Na 2 S0 4 , by heating with 
sulfuric acid in a shallow cast iron pan a mixture of two moles 
of salt with one mole of " chamber" sulfuric acid. A rapid 
reaction takes place until half of the salt has been decomposed 
forming sodium acid sulfate and hydrogen chloride as shown 
in the following equation: 

Nad + H 2 S0 4 = NaHS0 4 + HC1 

The evolution of hydrogen chloride now comes to an end and the 
mixture solidifies. This mixture is raked out of the pan on to 

FIG. 53. 

the hearth of a reverberatory furnace where it is kept well 
stirred by rakes and heated to a high temperature by being 
brought in contact with the flames and hot gases from the fire 
end of the furnace. Such a furnace together with the decom- 


posing pan is shown diagrammatically in Fig. 53, p is the decom- 
posing pan, h the hearth, f and f the fires and c, c the flues by 
which the hydrogen chloride and the products of combustion 

The furnace is called a reverberatory furnace because the 
flames are deflected by the roof and caused to play upon the 

On the hearth of this furnace the sodium acid sulfate reacts 
with the remaining sodium chloride as shown below, 

NaCl + NaHS0 4 = Na 2 S0 4 + HC1 

The sodium sulfate so obtained is called salt cake. 

The second step consists in the reduction of the sodium sulfate 
to sulfide, Na 2 S, by powdered coal; and the third, of the trans- 
formation of the sulfide to carbonate by heating with calcium 
carbonate in the form of chalk, or powdered limestone. 

These two reactions are carried out at one operation by heating 
a proper mixture of sodium sulfate, slack coal and limestone, 
either in a reverberatory or in a rotary cylindrical furnace through 
which the flames from the fire box pass. The latter form of 
furnace saves the great amount of hand labor required to keep 
the charge stirred on the hearth of the reverberatory furnace. 
The equations for the reactions are, 

Na 2 S0 4 + 2C = Na 2 S + 2CO 2 
Na 2 S + CaCOs = Na 2 C0 3 + CaS 

The product of this fusion is black in color and is known as 
black ash. It contains sodium carbonate, calcium sulfide, 
calcium oxide, coal, and a large number of other susbtances. 
Sodium carbonate is the only easily soluble substance present 
hence it is extracted from the black ash by water, using the 
principle of counter currents by allowing the fresh water to come 
in contact with the ash from which nearly all the sodium car- 
bonate has been dissolved while the nearly saturated solution 
acts upon the fresh black ash. In this way, practically all the 
carbonate is removed from the ash and a nearly saturated solu- 
tion is obtained. This solution is then evaporated and the salt 
separates as the monohydrate, Na 2 C03'H 2 O, which is called 
" crystal carbonate." The crystals are dried at a higher tern- 


perature forming the anhydrous salt, Na 2 C03, known as soda 
ash. By dissolving this in hot water and allowing it to crystallize 
at ordinary temperatures, the deca-hydrate, Na 2 C0 3 -10H 2 O, 
crystallizes out. This is known as sal soda, or washing soda. 
It dissolves more rapidly than soda ash and is largely used for 
domestic purposes, but contains so much water that it does not 
pay to ship it far. 

The sulfur in the calcium sulfide must be recovered for two 
reasons, first, competition is very strong in the soda industry 
and, second, if the sulfide were exposed to the air and weather it 
would be gradually dissolved and pollute the neighboring streams, 
and since under the action of the carbon dioxide of the air and 
water, hydrogen sulfide would be evolved, these streams would 
be very offensive. 

One of the best ways to recover this sulfur is to suspend the 
residue in water and pass in carbon dioxide from lime kilns when 
hydrogen sulfide is formed according to the following equation: 

CaS + H 2 + C0 2 = CaC0 3 + H 2 S 

This hydrogen sulfide is then burned to water and sulfur in a 
limited supply of air, 

2H 2 S + O 2 = 2H 2 O + 2S 

The hydrogen chloride evolved in preparing the salt cake is of 
course saved and either marketed as hydrochloric acid or made 
into bleaching powder. 

The Solvay or ammonia soda process has been in successful 
use for more than forty years. It consists essentially in saturat- 
ing a solution containing sodium chloride and ammonia in the 
proper proportions with carbon dioxide. When the solution 
becomes well saturated, solid sodium bicarbonate, NaHCOa 
begins to separate out and continues to do so until a little more 
than two-thirds of the sodium chloride has been transformed. 

While the mechanical details of the process are so complex 
that some of the larger works on technical chemistry should be 
consulted for them, its chemistry is simple. The carbon dioxide, 
ammonia, and water combine to form a solution of ammonium 
bicarbonate, NH 4 HC0 3 , 

C0 2 + NH 3 + H 2 O = NH 4 HCO 3 


This then reacts with the sodium chloride to form sodium bicar- 
bonate which is precipitated, and ammonium chloride which 
remains in solution. The equation is, 

NH 4 HC0 3 + NaCl<^ NaHCOs + NH 4 C1 
or written in the ionic form 

NH 4 + + HC0 3 " + Na + + CP^NaHCOs + NH 4 + + C 

The sodium bicarbonate is then filtered off, washed, and heated 
to change it into the carbonate, 

2NaHC0 3 = Na 2 C0 3 + C0 2 + H 2 O 

The carbon dioxide so obtained furnishes a portion of that re- 
quired in the process, but the greater part of the gas comes from 
the burning of limestone in specially designed lime kilns. The 
ammonium chloride which is left in the solution from which the 
sodium bicarbonate has separated is worth about eight times as 
much as the sodium carbonate which it helps to prepare, conse- 
quently the ammonia must be recovered and used again. This is 
done by adding slaked lime to the mother liquor and boiling off 
the ammonia formed, absorbing it in fresh brine and so keeping 
it in circulation. The equation is, 

2NH 4 C1 + Ca(OH) 2 = 2NH 3 + 2H 2 O + CaCl 2 

The calcium chloride so obtained is mixed with undecomposed 
sodium chloride and finds but little application. 

The Solvay process produces much purer sodium carbonate 
than the Le Blanc process and is for many locations the cheaper. 

Properties of Sodium Carbonate. The deca-hydrate forms 
large transparent crystals which effloresce easily, passing over 
into the monohydrate. This, in turn, in very dry air or at 
a higher temperature loses water and becomes anhydrous. 

Anhydrous sodium carbonate is a white opaque substance. 
It dissolves in water with the evolution of heat. A solution 
of sodium carbonate is fairly alkaline, due to hydrolysis since 
carbonic acid is a weak acid. For solutions of ordinary con- 
centration, something like 1 to 2 per cent, of the carbonate is 


Uses of Sodium Carbonate. Sodium carbonate is used in 
making glass, soap, in the softening of water, in washing and 
bleaching of linen and cotton fabrics, washing wool, paper mak- 
ing, dyeing and dye manufacture, the making of sodium salts, 
as a reagent in the laboratory to furnish the carbonate ion, and 
for many other important and useful purposes. Its consumption 
amounts to millions of tons yearly. 

Sodium Bicarbonate. Sodium bicarbonate, NaHCO 3 , which 
is also called sodium acid carbonate, sodium hydrogen carbonate, 
primary sodium carbonate, monosodium carbonate, baking soda, 
saleratus, or by the housewife simply soda, occurs in nature along 
with sodium carbonate. It is obtained in the manufacture of 
sodium carbonate by the Solvay process or by acting upon the 
monohydrate with carbon dioxide. 

It is a white crystalline powder which decomposes into sodium 
carbonate, water, and carbon dioxide, when heated to a moderate 
temperature. A solution of the bicarbonate contains a little 
carbonic acid formed according to the following equation. 

2NaHC0 3 ^ Na 2 C0 3 + H 2 CO 3 

When the solution is boiled, the carbonic acid breaks down into 
water and carbon dioxide, which escapes in the form of bubbles. 

Sodium Nitrate. Sodium nitrate, NaN0 3 , occurs widely 
distributed in nature in small amounts and is found in enormous 
quantities in certain deposits in the desert regions of Chile. 
These deposits are scattered over an area a few miles in width 
and 500 miles long. The crude nitrate always contains con- 
siderable sodium chloride, nitrate and perchlorate of potassium, 
and sodium iodate, with occasionally some chromate. The 
iodine is an important by-product. The nitrate is purified by 
dissolving it out of the crude material with hot water and allow- 
ing the solution to crystallize. 

The yearly production is about 2,000,000 tons which is some- 
thing like 1 per cent, of the available supply. 

Sodium nitrate crystallizes in rhombohedra which melt with- 
out decomposition at 317. It is very soluble in water and 
deliquesces, which makes it unfit for the better grades of gun- 
powder. It is used in the manufacture of blasting powder, 
of nitric and sulf uric acids, of potassium nitrate, as a preservative 



for meats, and in making sodium nitrite, but the greater part is 
used as a nitrogen fertilizer. 

Sodium Nitrite. Sodium nitrite, NaN0 2 , is made by heating 
sodium nitrate with lead or iron and recrystallizing the product, 

NaNOa + Pb = NaN0 2 + PbO 

It is very soluble in water though less so than the corresponding 
potassium salt. It is largely used in the manufacture of organic 

Sodium Sulfate. Sodium sulfate, deca-hydrate, Na 2 S0 4 -10H 2 
has been known for centuries under the name of Glauber's 
salts, having been first described by Glauber in 1658. The 
salt may also be obtained as the hepta-hydrate, Na 2 S0 4 -7H 2 O 
and in the anhydrous state. 

As has been mentioned, sodium sulfate is made as the first step 
in the Le Blanc process for the preparation of sodium carbonate 
and also for use in the manufacture of glass. 

The solubility relations between the different forms of sodium 
sulfate are interesting and instructive, being in many ways 
typical of all salts which crystallize with different numbers of 
molecules of water of crystallization. 

, Cms 
per 1C 


Na 2 SO t 
Grams Water 

















-10 0" 10 20 30 40' 50 60 70 80C 

FIG. 54. 

The two hydrates and the anhydrous salt are each distinct 
substances and as such, each has its own solubility at any given 
temperature, which, in general, is different from that of either of 
the others, as is shown by the curves in Fig. 54. 

These curves show, for example, that at 20 a solution in 


equilibrium with Na 2 S04-10H 2 will contain per 100 grm. of 
water 19.4 grm. of the salt calculated as Na 2 S04, while those in 
equilibrium with Na 2 S0 4 -7H 2 O or with the anhydrous salt con- 
tain respectively 44 grm. and 54 grm. of Na 2 S0 4 per 100 grm. of 
water. These figures demonstrate that the solubility of a salt 
cannot be expressed by a definite number unless there is given, 
in addition to the temperature and pressure at which the meas- 
urement is made, the particular solid phase with respect to which 
the solution is in equilibrium. 

Sodium sulfate is especially prone to form supersaturated 
solutions and, if a solution be prepared which is saturated or 
nearly saturated with the salt at about 32, but which does not 
contain the smallest trace of the solid deca-hydrate, be cooled 
to ordinary temperatures in a flask closed with a plug of cotton 
all the salt will remain in solution, although it contains about 
twice the solute which a saturated solution of the deca-hydrate 
would contain. If now even a very small piece of a crystal of 
Na 2 S04-10H 2 be introduced into the liquid, crystallization will 
at once take place and continue until the solution is no more than 
saturated. Sodium sulfate is such a common substance that it 
is almost universally present in dust and the cotton plug is used 
to exclude the latter, and so prevent the crystallization of the 

The solubility curves for the deca-hydrate and the anhydrous 
salt intersect at 32.4 and hence at this temperature the two 
salts are in equilibrium and have the same solubility. At all 
temperatures below this point the anhydrous salt is the more 
soluble and is meta-stable toward the deca-hydrate, while above 
this point the deca-hydrate is the more soluble and is meta-stable. 
In the curves, meta-stable equilibrium is shown by the dotted 
and the stable by the solid lines. 

When crystals of the deca-hydrate are heated to 32.4, they 
partially liquefy depositing some anhydrous salt and forming a 
solution saturated both with the deca-hydrate, and the anhydrous 
salt. The crystals of the deca-hydrate give a rather high vapor 
pressure of water and hence effloresce easily upon exposure to 
the air. 

Sodium sulfate is used in medicine as a cathartic and also in 
the manufacture of glass and of sodium carbonate. 


Sodium Sulfite. Sodium sulfite, Na 2 S0 3 -7H 2 0, is prepared by 
dividing a solution of sodium carbonate into two equal parts and 
saturating one with sulfur dioxide when carbon dioxide and acid 
sodium sulfite will be formed, 

Na 2 C0 3 -f H 2 + 2S0 2 = 2NaHS0 3 + C0 2 

The remainder of the carbonate solution is then added and carbon 
dioxide and the neutral sulfite will be produced, 

Na 2 C0 3 + 2NaHS0 3 = 2Na 2 S0 3 + H 2 + C0 2 

In addition to the acid and neutral salts mentioned above, a 
compound having the formula, Na 2 S 2 05, known as sodium disul- 
fite or sodium meta-bisulfite is made in rather large quantities 
for use in photography, as is also the neutral sulfite. 

When exposed to the air, the sulfites take up oxygen and are 
changed to sulfates. This process is generally retarded by the 
presence of small quantities of certain organic substances such 
as alcohol or sugar which here act as negative catalyzers. The 
use of the sulfites in photography depends largely upon their 
power of reacting with oxygen since they are added to the 
"developers," which are strong reducing agents to protect them 
from the oxidizing action of the air. 

Sodium sulfite is also used as an antiseptic or preservative and 
as an "antichlor" to reduce any chlorine which may be left in 
the goods after bleaching with this substance. 

Sodium Sulfide. Sodium sulfide, Na 2 S, is prepared on a large 
scale by the reduction of the sulfate with powdered coal at a 
high temperature as was done in the preparation of sodium car- 
bonate by the Le Blanc process. It dissolves in water and the 
solution has an alkaline reaction owing to 'hydrolysis which 
results in the formation of some of the hydrosulfide, NaHS, or 
rather of the hydrosulfide ion, HS~. The salt crystallizes with 
nine molecules of water as Na 2 S-9H 2 0, and is able to take up 
sulfur to form various polysulfides of which the compounds 
Na 4 S 9 -14H 2 0, and Na 2 S 5 are the best known. 

Small quantities of these sulfides are used in the laboratory as 
reagents, but the salts find extensive use in the manufacture of 
some kinds of glass and in the removal of hair from hides. 



Sodium Thiosulfate. Sodium thiosulfate, Na 2 S 2 3 -5H 2 0, is an 
important salt which may be made by boiling a solution of the 
sulfite with sulfur. The latter is taken up much as oxygen is by 
the sulfites, and thiosulfates are regarded as sulfates in which 
one atomic weight of sulfur has taken the place of an atomic 
weight of oxygen. The salt is obtained as a by-product of the 
Le Blanc process. The residue of calcium sulfide left in this 
process slowly oxidizes in the air to the thiosulfate, and when 
this is treated with sodium carbonate, calcium carbonate and 
sodium thiosulfate result. 

The penta-hydrate is very soluble and readily forms super- 
saturated solutions. It melts at 48 and if protected from dust 
will remain liquid at ordinary temperatures until brought in 
contact with some of the solid salt. This liquid may be regarded 
as either the supercooled hydrate or supersaturated solution. 

The salt is used as an "antichlor" to destroy the excess of 
chlorine used in bleaching and as a protection against the "war 
gas" which is chlorine, and also as a fixing agent in pho- 
tography, a use which will be discussed in connection with the 
silver salts. 

Oxyhalogen Compounds of Sodium. The methods of prepara- 
tion, properties, and uses of sodium hydrochlorite, NaCIO, 
and chlorate, NaClOs, have already been given under the oxygen 
compounds of chlorine and will not be repeated here. 

Sodium Phosphates. Three sodium salts of orthophosphoric 
acid are known, monosodium phosphate, NaH 2 PO4, disodium 
phosphate, Na 2 HP(>4, and trisodium phosphate, Na 3 PO4. The 
disodium salt, Na 2 HPO 4 -12H 2 O, is the most familiar of these. It 
is made by neutralizing a solution of phosphoric acid with sodium 
hydroxide or carbonate and allowing it to crystallize. The salt is 
easily soluble and effloresces readily. It is used in medicine as a 
laxative and in the laboratory as a soluble phosphate reagent to 
furnish the hydro phosphate HP0 4 , and the phosphate, 
PO4~" ", ions. Solutions of Na 2 HPC>4, are practically neutral in 
reaction, from which it follows that the ion HPO4 ~ ~ is but very 
slightly dissociated into the hydrogen and phosphate ions, so 
that the concentration of the latter is very small. For this 
reason, if the solution is to be used as a source of the phos- 
phate ion, P0 4 , sodium or ammonium hydroxide is added to 


reduce the concentration of the hydrogen ion and so make the 
HP04 ion break up. The reaction, . 

HPO 4 ^H++PO 4 

Ci C2 C 3 

is reversible and according to the law of mass action, C 3 c 2 /Ci = a 
constant. The addition of a base to the solution will decrease 
the concentration of the hydrogen ion and with it the product of 
02 and Cs and this will of course have to be followed by a decrease 
in Ci, the concentration of the hydrophosphate ion. But the 
hydrophosphate ion decreases in concentration by changing into 
the hydrogen ion, which is used up by the base, and the phos- 
phate ion which accumulates in the solution. So a solution of 
disodium phosphate to which sodium or ammonium hydroxide 
has been added contains a much larger amount of the phosphate 
ion than a pure solution of the salt and is used when the proper- 
ties of the phosphate ion are sought. 

The trisodium salt, Na 3 PO 4 -12H 2 0, is made by adding an ex- 
cess of sodium hydroxide to a solution of the disodium salt and 
evaporating to crystallization. Its solution in water is very 
alkaline and the salt is practically completely hydrolyzed into 
Na 2 HP04 and NaOH. It is used to soften or break "hard 
water." It does this by precipitating the calcium and magnesium 
of the water as phosphates. 

The monosodium salt is very soluble in water and the solution 
is acid toward litmus. 

The sodium salts of the other phosphoric acids are not of 
much importance, except perhaps the metaphosphate, NaPO 3 , 
which fuses easily and in this state dissolves many oxides and 
then solidifies to glass-like masses having characteristic colors. 
For this reason it is used in blowpipe analysis. 

Sodium Tetraborate. Borax or sodium tetraborate, Na 2 B 4 O 7 - 
10H 2 O, has been described under boric acid (see p. 261). Like 
the metaphosphate, when melted it dissolves metallic oxides, and 
gives characteristically colored glass-like substances. Because 
of this, it is used in blowpipe analysis. It is also used as a flux in 
hard soldering. 

Sodium Silicate. A salt having the composition Na 2 Si0 3 can 
be made by fusing together sodium carbonate and silicon dioxide. 


It is soluble in water and is called water glass. It is used in fire 
proofing wood, in making a kind of artificial stone or sand brick, 
being mixed with sand and lime for this purpose, and in the 
household for preserving eggs. 

Sodium Cyanide. Sodium cyanide, NaNC, is made on a large 
scale for use in the extraction of gold, for which purpose it has 
some advantages over potassium cyanide; since a given weight 
of sodium cyanide has in it a larger amount of cyanogen, the 
essential ion for the purpose, than the same weight of po- 
tassium cyanide. The salt is made by passing ammonia over 
metallic sodium heated to 300 to 400 in an iron vessel when 
sodium amide, NaNH 2 , is formed, 

2Na + 2NH 3 = 2NaNH 2 -f H 2 

This sodium amide is then brought in contact with charcoal 
heated to dull redness when sodium cyanide and hydrogen are 

NaNH 2 + C = NaNC + H 2 

Sodium cyanide is also made by fusing the mixture of calcium 
cyanamide and carbon obtained by heating calcium carbide, 
in an atmosphere of nitrogen with sodium carbonate. 

Na 2 C0 3 + CaCN 2 + C = 2NaNC + CaC0 3 

Like all other soluble cyanides, sodium cyanide is very 

Sodium Acetate. Sodium acetate, NaC 2 H 3 2 -3H 2 O, is a very 
easily soluble salt which is made by neutralizing acetic acid with 
sodium carbonate. It is used in the laboratory to decrease the 
acidity of a solution without making it neutral or alkaline. This 
is useful in analytical chemistry because of the fact that many 
precipitates will form in weakly acid solutions which are soluble 
in more highly acid solutions. To make the conditions favorable 
for the formation of such a precipitate in a solution which con- 
tains a strong acid, such as hydrochloric, all that is necessary is 
to add to the solution sodium acetate in sufficient quantity. 
The greater part of the hydrogen ion from the strong acid 
combines with the acetate ion to form undissociated acetic acid, 
because acetic acid is a weak acid and thus the acidity of the 
solution is decreased. 


Analytical Properties of Sodium. Like the salts of all other 
metals, the solutions of sodium salts show one set of properties in 
common, and this is really what is meant when it is said that they 
contain sodium as ion. In general, the statement of the ana- 
lytical properties of an ion consists in giving a list of the more 
common ions with which it will form difficultly soluble com- 
pounds. There is, however, no very difficultly soluble com- 
pound of sodium, so this element must be recognized in some other 
way. For this, advantage is taken of the fact that the colorless 
flame of the Bunsen burner is colored intensely yellow by very 
small quantities of sodium and when this flame is viewed through 
a spectroscope, two bright yellow lines are seen if the instrument 
is very powerful or the two may appear as one in the smaller 
instruments (see Frontispiece). This is an exceedingly delicate 
test and will detect 0.0000000003 grm. of sodium. Enough 
sodium is present in practically all substances to give this test 
and one must judge as to whether sodium is present as as incon- 
siderable trace or as an essential constituent from the length 
of time which the flame color lasts as well as its intensity. 

When the light from a powerful arc lamp is examined with a 
spectroscope a very bright continuous spectrum is seen. If the 
light from the arc passes through a Bunsen burner flame colored 
yellow by a sodium compound before it enters the spectroscope, 
black lines appear in exactly the position of the bright sodium 
lines when viewed in the absence of the arc light. In fact, the 
black lines may be made to change into bright ones simply by 
diminishing the intensity of the arc. This and other similar 
experiments lead to the conclusion that an incandescent gas will 
take up from a more intense source of light just that kind of light 
which it gives out, and when the more intense source of light is 
a solid, this will produce dark lines on a bright background. 
The dark lines come just where the incandescent gas alone would 
give bright lines. The solar spectrum consists of a bright back- 
ground crossed by a very great number of dark lines many of 
which are identical in position with bright lines given by terres- 
trial substances when in the incandescent gaseous state. From 
these facts it is supposed that the sun consists of an intensely 
heated core surrounded by an atmosphere of gases which while 
incandescent are much cooler than the core and which con- 


tain many of our well-known terrestrial elements. The presence 
of sodium for example, is clearly indicated by the fact that there 
are dark lines in the solar spectrum which exactly coincide with 
the bright lines of the sodium flame (see Frontispiece). 

In this way the presence of something like forty terrestrial 
elements in the sun has been ascertained. Among these may 
be mentioned hydrogen, helium (first discovered in the sun and 
then found in the earth), calcium, carbon, chromium, cobalt, 
iron, magnesium, maganese, nickel, silver, sodium, potassium, 
and vanadium. 


Occurrence. Potassium is never found free in nature, but 
occurs widely distributed in small quantities in the silicate rocks 
which compose the greater part of the earth's crust. On an 
average, potassium constitutes about 2.46 per cent, of the litho- 
sphere. This is the same as the sodium content of the rocks. 
The silicates when acted upon by water and carbon dioxide 
slowly decompose and their potassium and sodium content passes 
into solution. One would expect then that the rivers flowing 
down into the sea would carry approximately equal quantities 
of sodium and potassium salts, but this is not the case, since 
the sodium greatly predominates. The explanation for this is 
found in the fact that soils have the power to hold back the 
potassium salts by what is known as adsorption (see p. 228). 

Potassium salts are adsorbed by the soils while the sodium 
salts are not, and hence the latter easily find their way to the 
ocean. Because of this, the ocean contains relatively more 
of sodium salts than of potassium (see p. 274). 

In spite of this disproportion between the sodium and potas- 
sium content of the ocean, it might be expected that the condi- 
tions which would give rise to the great deposits of salt would 
produce smaller deposits of potassium compounds. But since 
the potassium content of the sea is so small in comparison with 
that of the sodium, the ancient seas would have had to become 
nearly completely dry before any potassium salts would be 
deposited. This would put such salts at the very top of the 
deposits where they would be extremely liable to be dissolved and 


washed away. So far as can be reasoned out from our present 
conditions, such deposits could only be formed when a great body 
of sea water dried up completely in a desert region and the 
deposits became covered so deep by dirt blown in from the sur- 
roundings that they were protected from any further action of 
water. At any rate the condition would be quite exceptional, 
and as a matter of fact only three large deposits of potassium 
salts are known; that in the neighborhood of Stassfurt, Germany, 
a smaller one at Kalusz in Austria, and another in northeastern 
Spain near Barcelona. The climate of these regions is now far 
from that of a desert, but there is some evidence that it was such 
at the time of the formation of the deposits. The discovery of 
beds of potassium salts in this country has frequently been 
reported of late, one of the most promising being that at Searles 
Lake in California and another the deposits of basic alum called 
alunite in Utah. 

Potassium Compounds and Plants. Growing plants seem to 
have an imperative need for compounds of the alkali metals, 
potassium and sodium. Those which grow on the land take up 
potassium compounds from the soil to such an extent that their 
ashes formerly constituted the chief source of these compounds. 
The plants growing in the sea and many of those along the sea- 
shore are rich in sodium salts. Some sea plants, however, 
notably the giant kelps of the Pacific, take up potassium com- 
pounds in such quantities that they may become an important 
factor in meeting the needs of this country for these very im- 
portant substances. Large plants for the commercial extraction 
of potash salts from kelp are now being put in operation. 

There is then a direct connection between the soluble potassium 
content of a soil and its fertility, for, of course, plants can only 
feed upon soluble substances. The soils usually contain large 
amounts of insoluble potassium compounds chiefly silicates, 
which slowly decompose under the action of the water and carbon 
dioxide, giving rise to soluble potassium salts and thus tending 
to keep up the supply of the latter in the soil. But if the land 
is kept under cultivation and the greater part of the crop removed 
as is often the case, soluble potassium compounds cannot be 
formed as rapidly as they are needed, and the fertility of the soil 
must decrease. The remedy is obviously to restore everything 


possible to the soil, and to make up the deficiency in potassium 
by use of some of the potassium salts as fertilizers. 

Animals have need for considerable quantities of sodium com- 
pounds chiefly for the chloride which constitutes the greater 
part of the salts in the body fluids, but their demand for potassium 
is rather small. Herbivorous animals get with their food much 
larger quantities of potassium and smaller quantities of sodium 
than they need. Now the elimination of the unnecessary po- 
tassium seems to be unavoidably accompanied by the excretion 
of the useful sodium compounds, and this accounts for the 
great appetite for sodium chloride shown by herbivorous animals. 
Carnivorous animals get their sodium and potassium in the proper 
proportions in their food and hence do not care especially for salt. 

Preparation of the Element. Metallic potassium was first 
prepared by Davy 1807. It is made by the same method that 
was used in the preparation of sodium, the electrolysis of the 
fused hydroxide. The demand for potassium is not large and 
only small quantities are made, because sodium which is much 
cheaper, will do nearly everything that potassium can do, and 
since the atomic weight of potassium is 39.1 while that of sodium 
is 23, 39.1 parts of potassium would be required to do the work 
which would be done by 23 parts of sodium at a much lower cost. 

Physical Properties. Potassium is a silver white metal which 
very quickly tarnishes upon exposure to moist air. It is waxy 
at ordinary temperatures, melts at 62.5 and boils at 758; 
the vapor is blue in color and. the molecular weight is apparently 
identical with the atomic weight, 39.1. Potassium easily forms 
an alloy with mercury which is much like the corresponding 
sodium amalgam. 

Sodium and potassium form an alloy which is interesting 
because, if it does not contain too much sodium, it is a liquid at 
ordinary temperatures. 

Chemical Properties. The chemical properties of potassium 
are practically the same as those of sodium except that potassium 
is somewhat more active, so that when thrown upon water the 
potassium gets hot enough to set fire to the hydrogen which is 
evolved and which burns with a violet flame, this being the color 
given to flames by potassium compounds. 

In fact, the entire chemistry of potassium is so similar to that 


of sodium that in most cases it will be necessary to add to what 
was given under sodium merely the slight points of difference 
presented by the potassium compounds. 

Potassium Hydride. Potassium hydride, KH, is prepared in 
the same way as the sodium hydride and is very similar to it in its 

Potassium Oxides. Potassium forms two compounds with 
oxygen, a monoxide, K 2 0, which dissolves in water to give 
potassium hydroxide and a peroxide, K 2 O 4 , which differs from 
the peroxide of sodium in composition, but like it, dissolves in 
water with the formation of the hydroxide, hydrogen peroxide 
and oxygen. 

Potassium Hydroxide. Potassium hydroxide, KOH, or caus- 
tic potash is so similar to sodium hydroxide that practically 
everything which was said concerning that compound might be 
repeated here as a description of potassium hydroxide. It may 
be prepared by the electrolysis of potassium chloride or by acting 
upon a dilute boiling solution of the carbonate with calcium 
hydroxide (milk of lime), calcium carbonate being precipitated 
and potassium hydroxide left in solution. It is exceedingly 

It is a very strong base and is used as such, but wherever possible 
sodium hydroxide is used in its place. The reason for this is 
that the sodium hydroxide is cheaper and less of it is required to 
do a given amount of chemical work because of the smaller 
atomic weight of sodium. 

Potassium hydroxide is very deliquescent, and its solutions 
take carbon dioxide from the air forming the carbonate which is 
also deliquescent. 

The chief technical use for potassium hydroxide is in the manu- 
facture of soft soap and of oxalic acid. 

Potassium. Chloride. Potassium chloride, KC1, is found in the 
potassium deposits of Stassfurt and is known as sylvite. It is 
also found there as double salts with magnesium chloride, 
KCl-MgCl 2 -6H 2 O, known as carnallite and with magnesium sul- 
fate, MgS0 4 -KCl-3H 2 O, kainite. The greater part of the potas- 
sium chloride of commerce is made from carnallite. The crude 
carnallite is crushed and heated with a solution of magnesium 
chloride left from previous operations. In this the carnallite 


easily dissolves. Upon cooling the solution the greater part of 
the potassium separates as the chloride, which is removed and 

Potassium chloride is easily soluble in water, and unlike sodium 
chloride, it increases markedly in solubility with rising tem- 
perature. It is the source from which most of the other po- 
tassium compounds are made. 

Potassium Bromide. When bromine is dissolved in potassium 
hydroxide solution, water and potassium bromide and bromate 
are formed. 

6KOH + 3Br 2 = 5KBr + KBr0 3 + 3H 2 O 

If this solution be evaporated to dryness and the residue gently 
heated, the bromate will decompose into bromide and oxygen. 
The greater part of the bromide of commerce is made by first 
preparing a solution of a bromide of iron by heating together 
water, bromine and iron filings and then precipitating the iron 
with a solution of potassium carbonate. By filtering and 
evaporating the solution, the pure salt may be obtained. 

Potassium bromide is used in medicine as a sedative. Since 
it is a habit-producing drug, it must be employed with caution. 
Sodium bromide is gradually replacing the potassium compound 
because the potassium ion is a dangerous heart depressant, 
while the sodium ion is if anything a heart stimulant. The 
sedative qualities of either drug are due to the bromine ion. 
Potassium bromide is also used in photography and in the labora- 
tory when a soluble bromide is needed as a source of bromine 
as ion. 

Potassium Iodide. Potassium iodide, KI, may be prepared by 
either of the methods used for the bromide. It is very soluble. 
Iodine is much more soluble in a solution of potassium iodide 
than in water forming the salt KI 3 which apparently gives the 
ions K+ + I 3 ~. 

Potassium iodide is used in medicine, in photography, and in 
the laboratory when a soluble iodide is wanted as a source of 
the iodine ion. 

Potassium Chlorate. Potassium chlorate, KC10 3 , is the most 
important of the oxyhalogen salts of potassium. The methods 


for its preparation, and its uses have already been discussed on 
pp. 134 and 135. 

Potassium Perchlorate. When potassium chlorate is cau- 
tiously heated, a part of it is converted into the perchlorate, 
KC1O 4 . This is a white crystalline substance which is not very 
soluble in water. Because of this, it is formed when perchloric 
acid is added to a solution of a potassium salt. It is almost 
insoluble in alcohol and therefore potassium is sometimes quan- 
titatively determined by precipitation as the perchlorate in 
alcoholic solution. As described on p. 139, the perchlorate 
decomposes at a high temperature into the chloride and oxygen. 

Potassium Nitrate. Potassium nitrate, KNO 3 , saltpeter, has 
been known for thousands of years being found as a white 
efflorescence in the neighborhood of human habitation in hot 
countries. The salt is formed by the action of oxygen of the 
air under the influence of nitrifying bacteria upon the nitrogenous 
materials and wood ashes, containing potassium carbonate, which 
accumulates under such conditions. India was the chief source 
of the salt for a long time. At the present day the greater part 
of the salt is made by bringing together equivalent quantities of 
sodium nitrate and potassium chloride, and a little water. The 
mixture is then heated to a fairly high temperature. Under 
these conditions a large amount of sodium chloride separates. 
The solution is then filtered while still hot and allowed to cool, 
when potassium nitrate and a little sodium chloride crystallizes 
out. The mixture is purified by washing with the mother liquors 
of previous recrystallizations, which removes the greater part of 
the sodium chloride, and then recrystallizing. 

This process may be easily understood after an inspection of 
Fig. 55, which gives the solubility curves of KN0 3 , KC1, NaCl, 
and NaN0 3 , the four possible salts resulting from the mixture 
of KC1 and NaN0 3 . 

It will be seen from these, that at high temperatures KN0 3 is 
the most soluble and NaCl the least so of the four salts, so in the 
hot solution NaCl separates. Upon cooling the solution, KN0 3 
becomes the least soluble and crystallizes out. 

Potassium nitrate is a white crystalline salt easily soluble in 
water with the absorption of heat, and having a bitter cooling 
taste. When heated to a high temperature, it decomposes, 



giving oxygen as one of its products. It was used by Scheele 
as one of the sources of this substance at the time of his discovery 
of the element. 

Potassium nitrate is used in the laboratory, in medicine, in 
preserving meats to which it imparts a red color, and in the manu- 
facture of fireworks and gunpowder. This last use absorbs by 
far the greater part of the KNOs of commerce. 

The composition of gunpowder varies somewhat but is near 
75 per cent, potassium nitrate, 10 per cent, sulfur, 14 per cent. 

1 25 

I 20 


40 60 

Degrees CerrHgrade 

Fit. 55. 



charcoal, and 1 per cent, water. The solid ingredients are pulver- 
ized separately, mixed together, moistened and ground until 
thoroughly incorporated. The mixture is then subjected to 
great pressure and afterward broken up into fragments which are 
sorted into sizes by sifting. The grains are polished or given 
a glaze by rattling them in a barrel, dried to remove the water 
added during the various operations, and finally dusted by 
passing the powder over very fine sieves. 

Gunpowder contains two easily combustible substances, sulfur 
and charcoal in intimate contact with a very good oxidizing 
agent, potassium nitrate. At a somewhat elevated temperature, 
these react very rapidly with the production of the gases nitro- 
gen, carbon monoxide, carbon dioxide, hydrogen, marsh gas, 
hydrogen sulfide and a number of solid substances chief among 
which is potassium carbonate. The great amount of gas pro- 
duced together with the high temperature of the reaction, about 
2,200C., produces great pressure which would reach above 6,400 


atmospheres or 47 tons per square inch if the products of the 
explosion occupied the volume of the, gunpowder before the 
explosion. To this high pressure the well-known effects of 
gunpowder are due. 

Sodium nitrate is much cheaper than the potassium salt, and 
is used in making the lower grades of powder for blasting, but 
cannot be used in the better kinds because it is deliquescent. 

Potassium Carbonate. Potassium carbonate, K 2 CO 3 , which is 
called potash or pearlash was formerly the most important of the 
potassium salts as it was the source of other potassium com- 
pounds, but since the discovery of the Stassfurt deposits it has 
largely lost its position. In former times, it was obtained from 
wood ashes by leaching them, using the principle of counter 
currents, but it is now chiefly obtained from the potassium sulfate 
of the Stassfurt deposits by a modification of the Le Blanc 
process for the preparation of soda; or from potassium chloride 
by taking advantage of the fact that when carbon dioxide under 
pressure is passed into potassium chloride solution containing 
magnesium carbonate in suspension, a double salt having the 
formula KHCO 3 -MgCO 3 -4H 2 O is deposited at temperatures 
below 24C. This salt is decomposed by hot water, forming 
MgCOs which is precipitated, carbon dioxide which passes out of 
the solution, and a solution of potassium carbonate from which 
the salt may be obtained. 

Potassium carbonate is also obtained from the residue of beet 
sugar manufacture, and some is still extracted from wood ashes. 

Potassium carbonate is a white solid; it is deliquescent and 
hence is very soluble in water. The solution is alkaline in; 
reaction because of hydrolysis and the formation of HC0 3 ~ + ; 

It is used in making soft soaps, hard glass, and for the prepara- 
tion of other potassium salts. 

Potassium Bicarbonate. Potassium bicarbonate, KHCO 3 , 
potassium acid carbonate or primary potassium carbonate as it 
is called, like the corresponding sodium compound, is much less 
soluble than the normal carbonate and may be prepared by 
passing carbon dioxide through a strong solution of the carbonate. 

Potassium Cyanide. Potassium cyanide, KNC, is an im- 
portant salt which is manufactured by heating potassium 


ferrocyanide, K 4 Fe(NC) 6 , with potassium carbonate when the 
following reaction takes place, 

K 4 Fe(NC) 6 + K 2 CO 3 = 5KNC + KNCO + CO 2 + Fe 

The salt KNCO is potassium cyanate and cannot be easily 
separated from the cyanide. If the ferrocyanide is heated alone, 
carbide of iron and nitrogen are formed besides the potassium 
cyanide. This process gives pure KNC but wastes cyanogen. 
By heating a mixture of dry potassium ferrocyanide with sodium, 
a mixture of potassium and sodium cyanides is obtained which 
is as good for most purposes as potassium cyanide and is sold as 
such. The equation is, 

K 4 Fe(NC) 6 + 2Na = 4KNC + 2NaNC + Fe 

Potassium cyanide is now made by the action of ammonia 
upon a fused mixture of potassium carbonate and charcoal. 
The reaction is apparently much like that for the preparation of 
sodium cyanide from the metal, ammonia and charcoal; for 
metallic potassium may be formed by the action of carbon on 
the carbonate. 

Potassium cyanide is very soluble in water and the solution is 
intensely and very rapidly poisonous. Hydrocyanic acid is a 
very weak acid, so solutions of potassium cyanide react alkaline 
owing to hydrolysis. The hydrocyanic acid formed can be 
readily detected by its odor. 

Potassium cyanide is used in large quantities in gold mining 
and in electroplating with gold and silver. It is used in the 
laboratory wherever a solution containing the cyanogen ion is 
required and also as a reducing agent. When it acts in the latter 
capacity, potassium cyanate KNCO is formed. When heated 
with sulfur, the latter is taken up forming a thio compound, 
analogous to the cyanates, known as potassium thiocyanate, 
KNCS, which is used in analytical work. This salt is sometimes, 
though incorrectly, called potassium sulphocyanide. 

Potassium Sulfate. Potassium sulfate occurs in nature in 
schoenite MgSO 4 -K 2 S0 4 -6H 2 O and several other Stassfurt salts. 
It is obtained from schoenite by adding potassium chloride and 
a little water, on heating, the comparatively difficultly soluble 


potassium sulfate crystallizes out and magnesium chloride is left 
in solution. 

It is used in the preparation of potassium carbonate and of 
alum, KA1(S0 4 ) 2 -12H 2 but chiefly as a fertilizer. 

Analytical Properties of Potassium. Solutions of potassium 
salt all have certain properties in common which are expressed 
by saying that they contain potassium as ion. 

Potassium compounds impart to the flame of the Bunsen burner a 
violet coloration which is easily masked by a little sodium unless 
the flame is viewed through a blue glass which cuts off the yellow 
sodium light and allows the violet of the potassium to be seen. 
The spectrum of potassium consists of a fairly strong line toward 
the extreme red end of the spectrum, a much fainter red line near 
the orange and a weak line well out in the violet end of the spec- 
trum. In between these lines is a host of faint ones so close 
together that in the smaller instruments they are not seen 
separately, but appear as a band of light from the orange to the 
violet (see Frontispiece). None of the potassium lines is any- 
where near as intense as the sodium line and hence the spectro- 
scopic test for potassium is far less delicate than that for sodium. 

Potassium in the ionic condition differs from sodium in that it 
will combine with a number of different ions to form compounds 
which are sufficiently difficultly soluble in water to make them 
useful for the detection and determination of this element 
Naturally for such a purpose the less soluble a salt the more 
useful it is, since the delicacy of the test and the accuracy of the 
determination depend directly upon this property. 

Potassium cobaltinitrite, K 3 Co(NO 2 )6, is the least soluble 
of the potassium salts, and is thrown down as a yellow precipi- 
tate when a solution of sodium cobaltinitrite, NasCoCNC^e, 
which contains the cobaltinitrite ion Co(N0 2 )6, is added to 
a neutral or slightly acid solution of a potassium salt, unless the 
latter be very dilute. This is a sensitive test for potassium, 
but is not as characteristic as the spectrum since ammonium, 
rubidium, and cesium compounds give very similar precipitates 
with the cobaltinitrite. But ammonium compounds are easily 
removed and rubidium and cesium occur so rarely that practically 
the test is a very good one for potassium. 

Potassium cobaltinitrite has a small but definite solubility 


which is diminished by the presence of an excess of either potas- 
sium or of cobaltinitrite as ion. In fact, it is a general rule that 
the solubility of a salt is diminished by increasing the concentra- 
tion of one of its component ions. 

Potassium as ion will also give rather difficultly soluble pre- 
cipitates with picric acid, CeH^NC^sOH, perchloric acid, 
HC1O4, hydrofluosilicic acid, H 2 SiFe, chloroplatinic acid, H 2 PtCl 6 , 
and tartaric acid, H 2 C 4 H 4 06, the precipitates being potassium 
picrate, CeH^NC^sOK, potassium perchlorate KC10 4 , potas- 
sium fluosilicate, K^SiFe, or potassium hydrogen tartrate 
KHC 4 H 4 06. The conditions which affect the solubility of this 
latter compound are so typical of those of other compounds that 
they are worthy of a careful study. 

When a strong solution of tartaric acid is added to a rather 
concentrated solution of a potassium salt, say the chloride or 
nitrate, the solution becomes supersaturated with respect to 
potassium hydrogen tartrate, KHC 4 H 4 Oe. This supersaturation 
may be so great that the solution will spontaneously deposit 
crystals, or it may persist until the solution is shaken vigorously 
or until a minute crystal of the salt is added. Of course after 
crystallization has once started it will continue until the solution 
becomes exactly saturated with the salt. After this point has 
been reached the addition of various reagents will produce inter- 
esting changes in the solubility of the substance. For example, 
the addition of a few drops of sodium hydroxide solution will 
cause a marked increase in the quantity of precipitate showing 
that the solubility of the potassium hydrogen tartrate has been 
decreased, but the addition of an excess of sodium hydroxide 
will cause the entire precipitate to disappear showing that an 
excess of this reagent increases the solubility of the potassium 

Further, the addition of either sodium acetate or of a very 
soluble potassium salt will decrease the solubility of the hydrogen 
tartrate, and cause a marked increase in the quantity of precipi- 
tate formed in the solution described above. 

The addition of a very moderate amount of one of the stronger 
acids such as hydrochloric, nitric, or sulfuric to a solution con- 
taining a precipitate of potassium hydrogen tartrate will increase 
the solubility of the latter and cause the precipitate to disappear. 


Acetic acid on the other hand has but little effect. Tartaric acid 
is a weak acid and this behavior of its potassium salt is typical 
of all difficultly soluble salts of weak acids, so that a general 
law may be put forward which says: The difficultly soluble 
salts of weak acids dissolve upon the addition of stronger acids. 
The converse of this is not true, the difficultly soluble salts of 
strong acids are not dissolved by weak acids nor are they, 
as a rule, dissolved by strong acids. 

Various more or less satisfactory explanations for these phe- 
nomena may be given, but the most useful and instructive is that 
offered by a combination of the theory of electrolytic dissocia- 
tion with the law of mass action. According to the theory, an 
electrolyte in solution is partially broken up into ions as shown 
below for acetic acid, 

Ci C2 Cs 

This is to be regarded as a reversible reaction and so subject to 
the law of mass action. Hence if Ci represents the concentration 
of the undissociated acetic acid and c 2 and c 3 that of the hydrogen 
and acetate ions respectively, equilibrium will result when, 

C 2 C 3 _ , 

If this be tested by experiment, the various concentrations being 
calculated from the electrical conductivity of the solution, the 
result will be found to agree closely with the theory as expressed 
in the mathematical relation just given. This is true not only 
for acetic acid but also for all other weak acids and for weak 
bases. It is true too, for dilute solutions of the more highly 
dissociated substances, the strong acids and bases and the salts, 
but apparently does not hold for the more concentrated solu- 
tions of these substances. This may be due to the conductivity 
being an imperfect measure of the concentration of the ions in 
such solutions, or to some other influence being superimposed 
upon those which are measured. Fairly satisfactory explana- 
tions of these deviations can be given, but are not of interest at 
this point because these deviations from the simple theory and 
law given above are not such as would alter any conclusion which 


we might reach as to the kind of change which would follow the 
addition of any given reagent, but would simply affect the calcu- 
lation as to the extent of such change. For qualitative purposes, 
then, the theory of electrolytic dissociation and the law of 
mass action may be combined and treated as though they exactly 
describe the actions of the substances. With this as preliminary, 
the discussion of the action of the various reagents upon potas- 
sium hydrogen tartrate as precipitated from solutions of potas- 
sium nitrate by solution of tartaric acid may be taken up in 
detail. The equation for the main reaction and equilibrium is as 
follows : 

K + + N0 3 ~ + H + + HC 4 H 4 6 - <^KHC 4 H 4 6 +H + +N0 3 ~ (I) 

Ci C2 'C 3 

The eqilibrium here is between the potassium ion, the hydrogen 
tartrate ion, and the undissociated potassium hydrogen tartrate. 
The concentration of these substances being represented respec- 
tively by Ci, 02, and c 3 , the mass law equation becomes, 


CiC 2 

It cannot be too strongly emphasized that c 3 represents the 
concentration of that portion of the potassium hydrogen tartrate 
which, while undissociated, is actually present in the solution in 
the dissolved state. It does not in any way represent that 
which is present as a precipitate. The only connection between 
c 3 and the precipitate is found in this, that when the solution is 
in equilibrium with some of the precipitate, i.e., is saturated, the 
equilibrium will be between the undissociated KHC 4 H 4 Oe in 
solution and that in the precipitate; and in this case, as in all 
others where there is equilibrium between the same substance 
in two phases, the equilibrium is determined by the ratio of the 
concentration of the substance in the two phases and not by the 
absolute amount of either phase. Therefore, since the concen- 
tration of the KHC 4 H 4 06 is fixed and constant in the solid phase, 
c 3 , that of this substance in solution must be fixed and constant 
also. From which it would follow that the product of Ci times 
02 would be constant for this saturated solution, and this point 
should be appreciated and kept thoroughly in mind. 


In addition to equilibrium I, given above, there are two others 
which are of importance in this connection. 

H + + HC 4 H 4 O 6 -^H 2 C 4 H 4 O6 (II) 

C 4 C 2 C 5 

The condition for equilibrium being, -^- = K' 



H + + C 4 H 4 6 - -^HC 4 H 4 6 - (III) 

C 4 C 6 C2 

Equilibrium resulting when, - - = K 

Now it must be remembered that all three of these equilibria are 
present at the same time in a solution which contains potassium 
nitrate and tartaric acid, and that anything which alters one 
changes each of the others. 

The addition of a few drops of sodium hydroxide to the solu- 
tion will affect first equation II by decreasing c 4 , the concentra- 
tion of the hydrogen ion, through the formation of water, but a 
decrease in c 4 will be followed by a decrease in c 5 and this will 
result in an increase in 02. But 02 comes in the equilibrium I 
and its increase at once produces an increase in c 3 . However, 
the solution was saturated with potassium hydrogen tartrate 
and hence the increase in c 3 makes it supersaturated and re- 
sults in the increase in the precipitate mentioned in the list of 
facts which were to be explained. 

The action of sodium acetate can now be easily understood 
since it acts like the hydroxide to decrease the concentration of 
the hydrogen as ion. 

An excess of sodium hydroxide, it will be recalled, redissolved 
the precipitate of potassium hydrogen tartrate; the explanation 
for this apparently contradictory behavior is as follows: As the 
hydroxide is added, water is formed decreasing c 4 and causing the 
undissociated tartaric acid to break up as shown in II until all the 
tartaric acid is transformed then the further addition of hydrox- 
ide will result in the breaking up of the hydrotartrate HC 4 H 4 Oe~ 
as shown in equation III; so by decreasing c 4 through the addi- 
tion of NaOH, 02 will be first increased and then decreased. But 


since c 2 is one of the factors in equation I, a decrease in c 2 will 
affect I, and will have to be followed by a decrease in c 3 . But if 
cs is decreased, the solution will become unsaturated and some 
of the precipitated substance will have to go into solution. So 
it may be easily seen why the solubility of potassium hydrogen 
tartrate is increased by the addition of an excess of sodium 

The increase in solubility of KE^H^Oe upon the addition of 
strong acids can be easily explained by considering the effect of 
the increase in the concentration of the hydrogen ion, c 4 , upon II. 
If 04 is increased, to keep the fraction c 5 /c 4 c 2 constant, c~ will have 
to increase which can only happen at the expense of c 2 , so c 2 will 
decrease. But in I a decrease in c 2 , as we have seen, is followed 
by a decrease in c 3 and this in turn by the solution of the precipi- 
tate. The failure of acetic acid to redissolve the precipitate 
can be understood, because acetic acid is so weak that it does 
not materially increase c 4 and therfore does not decrease c 2 and 
Cs, and hence leaves the solubility practically unaltered. 

While potassium hydrogen tartrate is a fairly important sub- 
stance, great quantities of it being used in the more expensive 
baking powders for example, it is not important enough to justify 
the time expended upon this discussion were it not for the fact 
that it is typical of most other substances, and that a thorough 
understanding of what is going on here will enable the student to 
deal easily with all other cases. 

Under sodium chloride it was mentioned that this salt would 
be precipitated from its saturated solution by the addition of 
hydrogen chloride. 

This may be understood by considering the following equili- 
brium : 

Na+ + Cl- * NaCl 

Ci C 2 C 3 

CiC 2 

If now hydrogen chloride be passed into the solution, c 2 will be 
increased, and this will of necessity be followed by an increase in 
Cs, and a decrease in Ci to correspond; but an increase in Cs makes 
the solution supersaturated, and brings about the precipitation of 
the salt. A strong solution of hydrochloric acid may be used 


instead of the gaseous hydrogen chloride with similar results. 
Potassium chloride will be precipitated in the same way and for 
the same reason by passing hydrogen chloride into its solution. 

As was mentioned in the preliminary to this discussion, page 305, 
this simple form of mass law treatment does not take into account 
all the factors and hence does not quantitatively describe the 
complex phenomena present in these cases, but it does fit them 
and others qualitatively. 


The methods of spectroscopic analysis were developed by 
Bunsen and Kirchoff about 1859 to 1860 and were immediately 
utilized by them in a search for new elements. In this they 
were successful for in 1860 they announced the discovery of a 
new alkali metal whose spectrum was characterized by two 
bright blue lines and hence was named Cesium, Cs, from caesius, 
the Latin word for the blue of the clear sky. Again, in 1861, they 
discovered another alkali metal whose spectrum consisted of two 
very strong lines in the violet and a number of others in the red, 
yellow and green portion of the spectrum. Two of the red lines 
are deeper red than the potassium lines, and hence the element 
was called Rubidium, Rb, from rubidius the Latin word for the 
darkest red color. 

Rubidium and cesium are always found in connection with 
lithium, sodium, and potassium. They are so like one another 
and so very much like potassium that it is a difficult matter to 
separate them not only from each 6ther but also from potassium. 

Metallic rubidium and cesium may be obtained by heating 
their corresponding hydroxide with powdered magnesium. 

Rubidium is a brilliant silver white metal which melts at 38.5 
and boils at 696. It has a density of 1.52 and acts upon water 
even more vigorously than potassium. The hydroxide so formed, 
RbOH, is at least as strong as that of potassium. The atomic 
weight of rubidium is 85.45. 

Metallic cesium has a silver white color, melts at 26.5 and 
boils at 670. Its density is 2.4. It is even more active chemic- 
ally than rubidium. Cesium hydroxide, CsOH, is a very strong 
base, and its salts are isomorphous with those of potassium and 
rubidium. The atomic weight of cesium is 132.81, 


Potassium, rubidium and cesium form a triplet of very closely 
related elements. The difference between the atomic weight 
of potassium, 39.1 and that of rubidium 85.45 is 46.35 which is 
just about that between rubidium and cesium which is 47.36. 
The same sort of relation exists between the triplet formed by 
chlorine, bromine, and iodine and here, too, the difference in 
atomic weight is about the same, being roughly 45. In a good 
many ways the properties of these three elements vary in a 
regular manner with the atomic weight as may be seen from the 
following table. 

As the atomic weight of the metal increases. 

The chemical activity of the metal increases. 

The solubility of the hydrotartrate increases. 

The solubility of the sulfate increases. 

The solubility of the chloride increases. 

The solubility of the chlorplatinate decreases. 

The solubility of the alum decreases. 

The boiling and melting points of the metals decrease. 


In the discussion of ammonia, NH 3 , it was pointed out that 
this substance would combine directly with acids forming 
ammonium salts. For example, ammonium chloride, NH 4 C1, 
is formed by the union of ammonia and hydrogen chloride, 

NH 3 + HC1 = NEUC1 

Ammonium chloride when dissolved acts as though it were 
dissociated into chlorine as ion and the complex cation NH 4 + , 
called the ammonium ion. This ion forms a long list of salts and 
is so very similar to potassium in its properties that it would be 
entirely out of place not to discuss it here in connection with this 
element. Most of the ammonium salts are isomorphous with 
the corresponding compounds of potassium, and the ammonium 
ion gives difficultly soluble compounds with the same reagents 
that the potassium ion does. So it is a rather easy matter to 
mistake the ammonium ion for the potassium. 

The sources and methods of preparation of some of the am- 
monium compounds have already been mentioned in connec- 


tion with the discussion of the compounds of nitrogen, pp. 

Ammonium. Hydroxide. Ammonium hydroxide is formed 
when ammonia dissolves in water, and may be obtained at low 
temperatures as a solid which melts at 79. Its solution 
shows comparatively weak basic properties indicating a small 
concentration of hydroxyl. 

Ammonium Chloride. Ammonium chloride or sal ammoniac, 
NH 4 C1, is one of the more important of the ammonium salts. 
It is easily soluble in water, and crystallizes in cubes or octahedra. 
Its solutions are slightly acid due to hydrolysis. When heated 
to about 350, it decomposes into ammonia and hydrogen chloride 
which pass off as gases, and recombine when they reach a cooler 
place, forming ammonium chloride, hence ammonium chloride 
is volatile. Most other ammonium salts undergo similar decom- 
position when heated giving ammonia and the acid. It is used 
in making dry cells and some other kinds of electrical batteries, 
as a reagent in the laboratory, in soldering and in medicine. 
Some idea of the extent of the use of ammonium chloride in 
batteries may be gathered from the fact that 50,000,000 dry cells 
are made in this country every year. 

Ammonium Nitrate. Ammonium nitrate, NH^NOs, is a white 
crystalline salt. At temperatures somewhat higher than its 
melting-point, 160, ammonium nitrate decomposes into nitrous 
oxide and water; 

NH 4 N0 3 = N 2 + 2H 2 

In this respect, it differs from most of the other ammonium 
salts since it does not yield ammonia and the acid. The differ- 
ence is presumably due to the oxidizing action of the nitric acid 
on ammonia. The nitrous oxide is a powerful oxidizing agent; 
for this reason the nitrate is used in some explosives, in fact under 
exceptional conditions it is explosive without any admixture. 

Ammonium Sulfate. Ammonium sulfate, (NH 4 )2SO 4 , is in 
some ways the most important of the ammonium salts. It is 
obtained in the recovery of ammonia in the gas works, and is the 
starting-point for the preparation of most of the ammonium 
compounds. It is extensively used for this purpose, but more 


as a nitrogen fertilizer, something like 500,000 tons of it being 
used for agricultural purposes each year. 

It is a white crystalline salt isomorphous with potassium 

The Sulfides of Ammonia. Ammonium hydrosulfide, NH 4 HS, 
is formed when equal volumes of ammonia and hydrogen sulfide 
are brought together. It is a crystalline substance which readily 
decomposes into its constituents. 

The salt is easily soluble in water and gives the ammonium 
and hydrosulfide ions. 

Two volumes of ammonia and one of hydrogen sulfide unite to 
form ammonium sulfide, (NH 4 ) 28. This is much less stable than 
the acid salt and passes quickly into the latter with the loss of 
amm6nia. When dissolved in water, it is largely hydrolyzed 
giving ammonium hydroxide and the hydrosulfide ion, HS~. 

Solutions of these salts may be made by passing hydrogen 
sulfide into ammonium hydroxide solution. They are used as 
reagents to furnish the sulfide ion S and hydrosulfide HS~~ 
which are necessary for the precipitation of such sulfides as 
required a higher concentration of sulfur as ion than is furnished 
by hydrogen sulfide, the sulfides of manganese and zinc for 

A solution of ammonium sulfide will dissolve sulfur, and 
thereby acquires a yellow color and is then known as yellow 
ammonium sulfide. It contains polysulfides and is used as a 

Ammonium Carbonate. When an excess of carbon dioxide is 
passed into a solution of ammonium hydroxide, ammonium 
bicarbonate, NH 4 HCO 3 , is formed. 

NH 4 OH + C0 2 = NH 4 HC0 3 

This is a fairly stable white crystalline salt which is soluble in 
water. The solid has a slight odor of ammonia and tends to 
decompose into ammonia, water and carbon dioxide. This 
reaction is quite rapid at moderately elevated temperatures and 
is employed by bakers who use the bicarbonate as a leavening 
agent much as the housewife does baking powder. For this 
purpose, the bicarbonate has the distinct advantage over the 


baking powder that it leaves practically no foreign substance in 
the food. 

An excess of ammonium hydroxide will convert the bicarbonate 
into the carbonate, (NH 4 ) 2 C0 3 , but this is very unstable and soon 

Commercial ammonium carbonate is made by subliming a 
mixture of the sulfate and powdered calcium carbonate. It is a 
mixture of the bicarbonate and carbamate, NH 4 C0 2 NH 2 . The 
latter is formed when ammonia combines with carbon dioxide. 

Ammonium Oxalate. Ammonium oxalate, (NH 4 ) 2 C 2 O 4 -2H 2 O 
is made by neutralizing oxalic acid with ammonium hydroxide. 
It is a white crystalline salt soluble in water and much used as a 

Ammonium Thiocyanate. Ammonium thiocyanate, NH 4 NCS ; 
is very easily soluble in water and is exceedingly poisonous. 
It is much used as a reagent. 

Analytical Reactions of Ammonium. The ammonium ion 
forms difficultly soluble compounds with the cobaltinitrite, chloro- 
platinate, and hydrotartrate ions. These compounds cannot be 
distinguished by the eye from the corresponding potassium 
compounds. Ammonium compounds, however, do not give a 
flame test or spectrum when introduced in the Bunsen flame. 

A very sensitive and characteristic test for ammonium com- 
pounds is based upon the fact that when a soluble base like 
sodium hydroxide is added to an ammonium salt, undissociated 
ammonium hydroxide is formed owing to the weakly basic 
character of this substance. As soon as the ammonium hydrox- 
ide is formed it breaks down into water and ammonia. The 
latter may be easily detected by its odor or by its action on 
moistened red litmus which is turned blue when held in the 
'vapors over the solution. The equations for reactions in the 
solution are as follows. 

NH 4 + + Cl- + Na+ + OH-^NH 4 OH + Na+ + C1-, 

NH 4 OH<=NH 3 + H 2 O 

The ammonia then vaporizes and dissolves in the water on the 
litmus where the following reactions take place 

NH 3 + H 2 0^NH 4 OH^NH 4 + + OH~ 
and it is this hydroxyl which turns the litmus blue. 


Of course potassium compounds can do nothing like this so it 
is very easy in this way to distinguish between the two sets of 

Ammonium Amalgam. The great similarity between the 
ammonium ion and the potassium ion suggests that if free am- 
monium, NH 4 , could be obtained it would be metallic in character. 
Attempts to prepare free ammonium have so far been unsuccess- 
ful, but an amalgam of this substance and mercury can be made 
either by electrolyzing most ammonium salts (the nitrate will 
not do), using a mercury cathode, or by the action of sodium 
amalgam on an ammonium salt solution* 

Below 85 the amalgam is a hard stable substance, but at 
ordinary temperatures it is very spongy and consists largely of a 
mass of bubbles due to the decomposition of the dissolved am- 
monium into hydrogen and ammonia. That ammonium has 
metallic characteristics is shown by the fact that it dissolves in 
mercury which will not dissolve anything else that is not a metal. 


The element lithium is very widely distributed and is obtained 
from petalite, LiAl(Si20 6 ) 2 , and lepidolite or lithium mica. The 
metal is prepared by the electrolysis of the fused chloride. It 
is a solid with a silver white luster, melts at 186 and boils above 
1400. Its density is 0.594 so that it floats on the petroleum used 
for its preservation. 

The atomic weight of lithium is 6.94 making it the third 
element in the order of atomic weights. It is active chemically 
but less so than the other members of the alkali group. It reacts 
rapidly with water but the metal does not melt and the hydrogen 
is not inflamed. It burns when heated in air, hydrogen, chlorine, 
bromine, iodine, sulfur, or nitrogen. This last property is inter- 
esting and somewhat important because there are not many 
things which will combine so readily with nitrogen. Lithium 
nitride, Li 3 N, is formed. 

Perhaps the most significant properties of lithium are that its 
phosphate, Li 3 PO4, and its carbonate, Li 2 COs, are rather diffi- 
cultly soluble in water, and that the bicarbonate LiHCOs is more 
soluble than the carbonate. With every other alkali metal, the 


bicarbonate is less soluble than the carbonate and the phosphate 
is easily soluble. The phosphates and carbonates of the members 
of the next group in the periodic system are difficultly soluble 
and the bicarbonates are more soluble than the corresponding 
carbonates. This shows that lithium has properties that make 
it a sort of connecting link or transition element between the 
two groups. That it really belongs to the alkali metals and not 
with the Group II is shown by the fact that it is monovalent while 
the members of Group II are divalent. 

Beryllium, the first member of Group II, is like lithium in that 
it has properties which connect it with both Groups II and III. 

In many ways it would have been more logical to have opened 
the discussion of the metals with that of lithium, but its post- 
ponement until this point just before taking up Group II is 
justified by the properties mentioned above. 

Lithium Carbonate. Lithium carbonate Li 2 CO 3 is commer- 
cially the most important lithium salt. It is used in medicine 
and for the preparation of lithium salts. 

Lithium Chloride. Lithium chloride is used in making artificial 
mineral waters and for making red fire and. other fireworks. 

Analytical Properties of Lithium. Lithium is very easily de- 
tected by the fact that it imparts a bright red color to the Bunsen 
flame and that its spectrum contains a very strong red line 
which is visible even when the lithium is present in exceedingly 
small quantities. 



Alkaline earth metals Magnesium sub-group 

General. The first member of Group II in the periodic system 
is beryllium. As mentioned under lithium, beryllium has proper- 
ties which connect it both with Group II and with Group III and 
therefore its discussion will be postponed until a later point. The 
other members of the group fall naturally into two well-marked 
sub-groups; the one called the alkaline earth metals and com- 
prising the common metals, calcium, strontium, and barium, and 
the very rare element radium; and the other sub-group called 
the magnesium family containing beryllium, magnesium, zinc, 
cadmium, and mercury. All the members of the group are 
divalent, but mercury also forms a series of compounds in which 
it is monovalent. 

Calcium, strontium, and barium form a triplet of very closely 
related elements, and for a long time have been grouped together 
under the name of the alkaline earth metals. The discussion 
of the group will be opened with that of these metals since they 
are in many ways more closely connected with the alkali metals 
than are the members of the magnesium sub-group. 


The compounds of calcium have been known and used from the 
very earliest times, but the element itself was first prepared, 
although in an impure state, by Davy in 1808. 

Occurrence. Calcium compounds are very abundant and are 
widely distributed over the earth's surface in the form of silicates, 
carbonates (limestone), sulfates, fluoride, phosphates, borates, 



etc. The element is present in the sun and certain of the fixed 
stars as shown by their spectra. 

Compounds of calcium are indispensable for the growth and 
life of plants and animals. The bones and teeth of animals are 
composed largely of calcium phosphate but contain some car- 
bonate and fluoride in addition. 

Preparation of the Element. Calcium is prepared on a 
fairly large scale by the electrolysis of the fused chloride, the 
anode consisting of carbon and the cathode of an iron rod. 
The calcium, as it is liberated, clings to the end of the rod and 
the latter is gradually withdrawn from the fused salt, causing 
the calcium to build up an irregular cylinder of the metal. It is 
essential that the temperature of the fused calcium chloride be 
kept as low as possible; for this purpose, calcium fluoride is 
sometimes added to the bath, since it dissolves and lowers the 
melting-point of the chloride. 

Properties. Calcium is a silver white crystalline metal which 
melts at 800 and volatilizes below its melting-point in a vacuum. 
It has a density of 1.55 and is malleable, although somewhat 
harder than lead. It tarnishes rapidly in moist air at ordinary 
temperatures, but may be turned and polished in the air if kept 

Calcium decomposes water, forming hydrogen and calcium 
hydroxide, Ca (OH) 2 . It acts vigorously on acids. When heated 
in the air, it burns brilliantly and also combines rapidly at higher 
temperatures with nitrogen, chlorine, hydrogen, sulfur, silicon, 
and phosphorus. With nitrogen,- calcium nitride, Ca 3 N2, is 
formed. When calcium is heated in a closed vessel, both the 
oxygen and nitrogen are removed and a very high vacuum 

The atomic weight of calcium is 40.07. 

Calcium Hydride. Calcium combines directly with hydrogen 
to form calcium hydride, CaH 2 , a white crystalline substance 
which reacts with water, forming the hydroxide and hydrogen. 

Calcium Oxide. Calcium oxide, or lime, is a very important 
substance as everyone knows. It does not occur in nature, but 
is readily prepared by heating marble or limestone to a high 
temperature when the following reaction takes place, 
CaCO 3 ?=* CaO + CO 2 





This reaction is reversible and the conditions governing it 
will be considered in detail under calcium carbonate. 

This process is carried out in furnaces called lime kilns. The 
more modern types of these kilns, Fig. 56, consist of a tall shaft 
of brick work with the interior somewhat egg-shaped in cross 

section, with two or more fire boxes 
located at the sides and opening into 
the shaft a short distance from the 
bottom. The flame and hot gases 
from the fuel are drawn into the shaft 
and heat the limestone above its 
temperature of decomposition, while 
the air drawn in by the draft of the 
stack sweeps away the carbon dioxide 
and so lowers the temperature neces- 
sary for the decomposition. The 
operation is continuous, the limestone 
being fed in at the top as fast as the 
lime is removed at the bottom. This 
country produces nearly four million 
tons of lime per year. 

Calcium oxide is a white substance 
having a density of 3.3. It shows 
some signs of melting when heated in 
the oxyhydrogen blowpipe and under 
such conditions gives a very bright 
light which is known as the calcium 
or lime light. 

The melting-point of lime is about 1900 and it not only melts 
easily at the temperature of the electric furnace, but even boils. 
When the vapors of the oxide condense, crystals of calcium oxide 
are formed. Calcium oxide is not reduced by sodium or by 
carbon except at temperatures of the electric furnace. The 
calcium obtained by the Deduction with carbon unites with 
more carbon to form the carbide, CaC 2 . 

Lime has a great tendency to combine with water forming the 
hydroxide, Ca(OH) 2 . Because of this, it is used in drying am- 
monia and certain other gases and in dehydrating alcohol. Its 

FIG. 56. 


most important use, however, is in the preparation of mortar and 

Calcium Hydroxide. Calcium hydroxide or slaked lime, 
Ca(OH) 2 , is made by the action of water on calcium oxide or 
quicklime as it is sometimes called. The equation is, 

CaO + H 2 O = Ca(OH) 2 > . 

The action is slow in starting but evolves considerable heat, 
so that after the reaction has once begun the temperature rises 
and it then goes on rapidly. When large quantities of lime 
slake in contact with wood, the temperature often rises 
high enough to set fire to the wood. In this way "buildings 
containing lime in storage are sometimes set on fire by water 
in times of floods. 

Calcium hydroxide is a white powder whose density is 2.08 
and which therefore occupies a larger volume than the lime from 
which it was formed. It is slightly soluble in water, 600 parts 
of the latter being required for one of the hydroxide at ordinary 
temperatures, and about twice as much water at the boiling- 
point since the solubility of the hydroxide decreases with rising 
temperature. The solution is decidedly alkaline in character 
because calcium hydroxide is a strong base, although not as strong 
as sodium hydroxide. The solution of the hydroxide is called 
lime water and is used in medicine and as a reagent for the 
detection of ca.rbon dioxide. 

Because of its cheapness, calcium hydroxide is largely used 
wherever a base is desired, unless the properties of the calcium 
ion interfere. If the presence of much water is objectionable, a 
suspension of the hydroxide in water called milk of lime is used. 
That part which is actually in solution reacts first and then more 
dissolves so that ultimately it may all be used. ' 

Slaked lime is used in the manufacture of sodium and potas- 
sium hydroxides, ammonia, bleaching powder, mortar, the 
softening of water, the purification of illuminating gas, the 
removal of hair from hides, and for many other purposes. 

Air-slaked lime is lime which has been exposed to the air 
until it has fallen to powder from the water and carbon dioxide 
taken up from the air; it contains calcium oxide, hydroxide, 
and carbonate. 


Mortar. Mortar is made by mixing together into a paste 
slaked lime, sand and water. The water gradually evaporates and 
the mixture sets. In this setting process, the calcium hydroxide 
decreases greatly in volume and renders the mass porous. The 
sand is added largely to decrease the shrink and to make the 
mortar more porous. After setting, the mortar gradually 
absorbs carbon* dioxide from the air and is converted into 
calcium carbonate, 

Ca(OH) 2 + CO 2 = CaC0 3 + H 2 O 

The carbonate crystallizes as it is formed and the crystals 
interlace and bind the sand and the masonry together. When 
this hardening process is complete the mortar is very hard and 
strong, but it takes place so very slowly that years and even 
centuries are required for its completion. The hardening does 
not take place under water nor in very damp places. Under 
such conditions hydraulic cement must be used. Such cements 
will be discussed under aluminum. 

Calcium Carbonate. Calcium carbonate occurs very abun- 
dantly in nature as limestone, marble, and chalk. The latter is 
made up of the fragments of the shells of marine organisms, while 
limestone and marble are composed of interlaced crystals of a 
form of calcium carbonate known as calcite. The only difference 
between the two is that marble is the more distinctly crystalline. 
Calcite is sometimes found in very large clear crystals and is then 
known as Iceland spar. There is another less stable crystalline 
form of calcium carbonate known as aragonite. Calcium car- 
bonate is difficultly soluble in water and is precipitated whenever 
a soluble carbonate is added to a solution of a calcium salt. 

The freshly precipitated carbonate is amorphous, and like all 
such substances is more soluble than the crystalline modifica- 
tion into which it passes after a time. At ordinary temperatures 
the change takes place slowly and calcite is formed, but at the 
boiling-point the transformation is rapid and aragonite is pro- 
duced. The solubility of calcite is so slight that a liter of water 
will dissolve only 12 mgrm. of the salt, but the solubility is 
greatly increased by carbon dioxide dissolved in the water owing 
evidently to the formation of the bicarbonate, Ca(HCO3)2, 

CaC0 3 


Underground waters are more or less heavily charged with 
carbonic acid, and hence dissolve some of the calcium carbonate 
with which they come in contact. 

Calcium bicarbonate is quite unstable and exists in solutions 
only when they contain an excess of carbonic acid. Therefore 
it is decomposed and the carbonate is precipitated when its 
solution is boiled or exposed to air, since in either event 
the greater part of the carbonic acid will be driven out of the 

The precipitation on boiling is the cause of the formation of a 
part of the scale which collects in boilers and in tea kettles, and 
the precipitation on exposure to air is the main cause of the 
formation of. stalactites on the roofs of caves and of stalagmites 
on the floors where the underground water drips down from the 
roof upon them. 

Waters containing calcium and magnesium salts are called 
hard waters, and that part of the hardness which is due to cal- 
cium bicarbonate is called temporary hardness because it may 
be removed by boiling for a time. It may also be removed by 
adding to the solution the proper amount of calcium hydroxide 
to change the bicarbonate into carbonate, 

Ca(HCO 3 ) 2 + Ca(OH) 2 = 2CaC0 3 + 2H 2 O 

The rest of the hardness is called permanent hardness, although 
it too may be overcome by proper chemical treatment as will be 
described in connection with the discussion of magnesium. 

As mentioned under calcium oxide the carbonate decom- 
poses at a high temperature into the oxide and carbon dioxide. 
The equation for the reaction is, 

CO 2 

Ci C2 C 3 

C_2CS = K 

This reaction is reversible, and applying the law of mass action 
the conclusion is reached that as C2C 3 /Ci = constant and since 
CaCOs and CaO are solids, Ci and c 2 are constant, that there- 


fore Cs the concentration of the CC>2 is constant for equilibrium 
at any one temperature. The concentration of a gaseous sub- 
stance is proportional to its pressure, so it follows that the carbon 
dioxide in equilibrium- with the CaO and CaC0 3 at any given 
temperature will always have a certain fixed pressure. 

The decomposition of the calcium carbonate is accompanied 
by the absorption of heat, and therefore in accordance with the 
law of mobile equilibrium the pressure of the carbon dioxide 
increases with rising temperature. It reaches atmospheric pres- 
sure at about 900, but in a stream of air or any other gas which 
will keep the partial pressure of the carbon dioxide low by sweep- 
ing it away as fast as it is formed, the decomposition will take 
place at a much lower temperature. In the lime kiln the lime- 
stone is heated in a current of gases in which the partial pressure 
of carbon dioxide is far lower than one atmosphere, and conse- 
quently the temperature of decomposition is comparatively low. 

Calcium carbonate although so slightly soluble in water 
readily dissolves in acids. This of course is another case of a 
difficultly soluble salt of a weak acid being dissolved by a stronger 

Calcium Chloride. Calcium chloride, CaCU, occurs in nature 
in sea water and as a constituent of a few minerals for example, 
tachhydrite, CaCl 2 -MgCl 2 -12H 2 0, and apatite, Ca 5 (PO 4 ) 3 Cl. 

It may be prepared by dissolving pure calcium carbonate in 
hydrochloric acid. It is a by-product of several technical proc- 
esses and is comparatively cheap. It is exceedingly soluble 
in water and crystallizes with varying amounts of water of 
crystallization forming a mono, di, two tetrahydrates and 
hexahydrate. The hexahydrate crystallizes from highly con- 
centrated solutions at ordinary temperatures. It is very deli- 
quescent and melts at 30. When mixed with ice in the best 
proportions, a temperature of 55 may be obtained. Anhy- 
drous CaCl2 and ice do not make a good freezing mixture since 
the salt evolves much heat in dissolving, while the hexahydrates 
absorb heat. 

When the hexahydrate is heated to 200, it loses four moles of 
water and forms a white porous mass of the dihydrate which is 
very hygroscopic and is used for drying gases. When heated 
to a still higher temperature the anhydrous salt is obtained but 


the reaction represented by the following equation takes place to 
a certain extent, 

CaCl 2 + H 2 = CaO + 2HC1 

and the product contains some calcium oxide. 

Calcium chloride forms a compound, CaCl 2 -8NH 3 , with 
ammonia and cannot be used for drying this substance. 

Calcium Hypochlorite and Bleaching Powder. Bleaching 
powder, CaCl 2 O, or chloride of lime, is made by the reaction of 
chlorine on slacked lime, 

Ca(OH) 2 + Cl a = CaCl 2 O + H 2 O 

In the most recent plants, the slacked lime is carried in a 
conveyer in a direction contrary to that of the chlorine so that 
the almost saturated lime meets the strongest chlorine, while 
the fresh lime comes in contact with the weakest chlorine. In 
this way, each is utilized to the greatest advantage. (For the 
properties and use of bleaching powder see pp. 128, 133.) 

Pure calcium hypochlorite, Ca(OCl) 2 , may be prepared by 
neutralizing hypochlorous acid with calcium hydroxide. It 
finds no application. 

Calcium Bromide and Iodide. Calcium bromide and iodide are 
very similar to the chloride, but are even more soluble than the 

Calcium Fluoride. Calcium fluoride, CaF 2 , or fluorspar as it 
is commonly called, occurs in nature fairly abundantly. It 
crystallizes in cubes and is nearly insoluble in water. It melts 
at 1,330 and has to a marked degree the property of reducing 
the melting-point of the slags produced in metallurgical processes 
and has long been used for this purpose. When heated it be- 
comes luminous at temperatures far below redness, and this 
gives rise to the term fluorescence. 

Calcium fluoride is the source of other fluorine compounds, 
and is used in the preparation of hydrofluoric acid for etching 
glass and for other purposes. 

Calcium Nitrate. Calcium nitrate, Ca(N0 3 ) 2 , is found in all 
fertile soils-. It may be made by the action of nitric acid on 
the carbonate, and is manufactured on a large scale for use as 
a fertilizer at the plant for the manufacture of nitrogen com- 


pounds from the air at Notodden, Norway. It is very soluble 
in water and forms a number of hydrates; that stable at ordinary 
temperatures is Ca(NO3)2'4H 2 O. The anhydrous salt is used 
as a drying agent. When heated it gives calcium oxide, nitrogen 
peroxide, and oxygen. 

Calcium Sulfate. Calcium sulfate, CaS0 4 , occurs in nature in 
two forms as anhydrite, CaSO 4 , and as gypsum, CaSO 4 2H 2 O; 
the latter is the more important since from it is made the valuable 
plaster of Paris and cement plaster. 

Calcium and sulfate ions are present in sea water and gypsum 
is deposited soon after the concentration of such water begins. 
Without doubt this has been the origin of the great beds of 
gypsum which are found in various parts of the world. The 
United States has many such deposits, notably in New York, 
Michigan, Iowa, Ohio, Texas, Oklahoma, and Kansas. The 
total production of gypsum in the United States is about 2,500,000 
tons per year, of which two-thirds is converted into plaster. 

Besides the dihydrate or gypsum, CaS0 4 -2H 2 O, calcium sulfate 
forms a hemihydrate, 2CaSO 4 -H 2 0, which is in equilibrium with 
the dihydrate and water vapor at 107 under a pressure of one 
atmosphere. When ground gypsum is heated from 110 to 130 
its aqueous tension is greater than one atmosphere, and it rapidly 
decomposes, losing water and changing into the hemihydrate. 
The product obtained in this way is called plaster of Paris. 
When mixed with water to form a paste, it quickly hardens to 
a mass of interlaced crystals of the dihydrate. In so doing it 
increases materially in volume and fills perfectly a mould into 
which it may be poured ; because of this it is much used in making 
plaster casts. By far the more extensive use for this product is 
as a plaster for the interior of buildings, and for such purposes 
the plaster of Paris sets too quickly. To overcome this, ad- 
vantage is taken of the fact that many organic substances such 
as glue, sugar, glycerine, etc., will greatly increase the time of 
setting; such substances are called retarders and the plaster so 
treated is known as cement plaster. Cement plaster may also 
be made without retarder by using gypsum which is mixed with 
the proper amount of clay or dirt. 

These plasters have many advantages over lime plasters for 
inside work and have very largely replaced them. One impor- 


tant advantage is that they dry quicker and are much harder 
and stronger than the lime plasters. If plaster of Paris is mixed 
with a solution of common salt instead of pure water, its time of 
setting will be shortened. 

Great care must be taken not to overheat the plaster or drive 
off all the water, since in these cases the plaster either will be 
too slow in setting or will not set at all. Such plaster is said to 
be "dead burned." 

The solubility of gypsum varies in a rather remarkable way 
with the temperature; from to 38, the solubility increases and 
then decreases as the temperature rises. At the boiling-point 
of water, in a steam boiler working under the usual pressure, 
calcium sulfate is nearly insoluble; and under such conditions, 
it is precipitated from water which contains much of this sub- 
stance, in the form of a scale which clings closely to the boiler. 
Hard boiler scale often consists largely of calcium sulfate. 
Waters containing this substance are said to have permanent 
hardness in distinction from temporary hardness due to the 
bicarbonate. The addition of the proper amount of sodium car- 
bonate will precipitate the calcium as carbonate and leave a 
solution of sodium sulfate. In this way the permanent hardness 
may be removed. 

Calcium Sulfide. Calcium sulfide, CaS, is made by heating the 
sulfate with charcoal. It is a by-product of the Le Blanc soda 
process, and is insoluble in water; but is slowly hydrolyzed forming 
calcium hydroxide and the hydrosulfide, 

2CaS + 2H 2 O = Ca(OH) 2 + Ca(HS) 2 

Many specimens of calcium sulfide especially if slightly impure 
with compounds of vanadium and bismuth will phosphoresce, 
i.e., shine in the dark, after being exposed to the light. The pure 
sulfide does not do this. 

Calcium Phosphate. The principal forms in which calcium 
phosphate occur in nature are apatite, Ca 5 (P0 4 )3F and the normal 
phosphate, Ca 3 (P0 4 )2. Its chief commercial source is from phos- 
pherite whose composition lies between that of apatite and the 
normal salt. Enormous beds of phosphate are found in many 
parts of the world; notably in this country in Florida, Tenn- 


essee, South Carolina, Utah and its neighboring States. The 
western deposit is the largest in the world and will doubtless at 
a future time furnish a large part of the supply. 

Calcium phosphate is present in nearly all rocks and all fertile 
soils. It is practically insoluble in pure water, but is very slowly 
dissolved by the carbonic acid which is present in the ground 
waters. Compounds of phosphorus are absolutely essential to 
all forms of life. Animals derive what they need principally 
from plants, while the latter get theirs from the phosphate 
rendered soluble by the carbonic acid. 

Although most soils contain in the aggregate. large quantities 
of calcium phosphate, the action of the weather in rendering 
this soluble is too slow to meet the demands of the crops under 
intensive cultivation and hence the plants must be fed with ready 
prepared soluble phosphates. For this purpose many millions 
of tons of phosphorite are mined each year and made into ' 'super- 
phosphate" fertilizer. 

The phosphate fertilizers are made by mixing approximately 
equal quantities of phosphorite and chamber-sulfuric acid. The 
dihydrates of calcium sulfate and monocalcium phosphate are 
formed as shown in the following equation : 

Ca 3 (P0 4 )2 + 2H 2 S0 4 + 6H 2 = Ca(H 2 P0 4 ) 2 -2H 2 O + 
2CaSO 4 -2H 2 O 

This reaction takes place with the evolution of heat, and the 
mixture soon solidifies to a mass composed of the two hydrated 
salts. This is known as superphosphate. The monocalcium 
phosphate is easily soluble in water, and is the more valuable 
part of the fertilizer. 

Besides the tricalcium phosphate or normal phosphate as it 
is called, Ca 3 (PO 4 ) 2 , there is known the dicalcium salt, CaHP0 4 , 
and the monosalt given above. All the calcium phosphates are 
readily soluble in acids, even in those as weak as acetic acid. 
This is due to the fact that phosphoric acid is weak and that ion, 
H 2 P0 4 ~, is formed with especial ease. The calcium salt of this 
ion is easily soluble as has been mentioned above. 

Besides being made into fertilizers, calcium phosphate serves 
as the source of phosphoric acid and phosphorus. 


Calcium Carbide and Calcium Cyanamide. Calcium carbide, 
CaC2, is formed by the reaction "of carbon upon lime in an 
electric furnace (Fig. 57). The equation for the reaction is 

CaO + 3C = CaC 2 + CO 

The carbide reacts with water producing acetylene, and with 
nitrogen for the formation of calcium cyanamide, CaCN 2 , 

CaC 2 + N a = CaCN 2 + C 

FIG. 57. 

This reaction takes place at 1,000 and goes more readily in the 
presence of calcium chloride. Calcium cyanamide is slowly 
decomposed by water, forming the carbonate and ammonia, 

CaCN 2 + 3H 2 = CaCO 3 + 2NH 3 

It is therefore useful as a nitrogen fertilizer, and is now being 
made at Niagara Falls in the United States, and in about fourteen 
different places in the world. While the nitrogen fertilizers are 
very important, they are not as absolutely indispensable as the 
phosphates and potassium fertilizers are since it is practicable 
to introduce the nitrogen compounds into the soil with leguminous 

Calcium Oxalate. Calcium oxalate, CaC 2 4 , is a white crystal- 
line salt which is very slightly soluble in water and is formed at 
once whenever a solution of an oxalate is added to a calcium 
salt. The precipitate is practically insoluble in water and in 
acetic acid, but dissolves easily in strong acids like hydrochloric 
and nitric. The explanation for the difference is to be found in 
the fact that oxalic acid is a stronger acid than acetic, but 
weaker than hydrochloric. It will be recalled that the difficultly 


soluble salts of weak acids are soluble in the stronger acids because 
of the decrease in the concentration of their anion due to the 
formation of the undissociated weak acid. The slight solubility 
of calcium oxalate is made use of in the detection and estima- 
tion of calcium as ion. The oxalate reagent used is ammonium 

Calcium Silicate and Glass. Calcium silicate occurs in nature 
in almost all of the natural silicates. The meta silicate, CaSiOs, 
forms the mineral wollastonite, but is not of much importance. 

The artificial complex silicates known as glass are, however, 
of great importance. The most common of these, that known 
as soft or soda glass, is made from calcium carbonate, sodium 
carbonate, and glass sand, which is nearly pure silicon dioxide. 
The reaction may be represented by the following equation: 

Na 2 C0 3 + CaC0 3 + 6SiO 2 = Na 2 CaSi 6 Oi 4 + 2CO 2 

but it must be understood that the glass is a mixture of silicates 
rather than a definite chemical compound. 

The substances are mixed in the proper proportions and heated 
to a temperature below their melting point when the reaction 
takes place, and most of the carbon dioxide escapes through the 
porous mass. The temperautre is then raised to about 1,200 
when the glass fuses to a mobile liquid which is fluid enough to 
allow the bubbles of gas to escape. The temperature is then 
lowered to 700 or 800 when the glass becomes viscous enough to 
work. Ordinary window glass is made by blowing long cylindrical 
bulbs of glass. After the bulb is blown to the proper dimensions 
the bottom is heated and blown out leaving a long cylinder with 
a bottle-like top and neck. This is allowed to cool and the top 
cut off by wrapping a string of hot glass around the cylinder near 
the top and suddenly touching a heated spot on the cylinder with 
a cold iron. A crack instantly extends around the cylinder 
following the line heated by the string of hot glass. The result 
is a long glass cylinder open at each end. This is then split 
lengthwise by running a red hot iron through the cylinder and 
touching the heated glass at the end with a cold iron. The 
crack instantly runs the whole length of the cylinder. This is 
then placed upon a flat revolving bed of a furnace upon which 
it is smoothed out after being heated to the softening point. 


Bottles are blown into moulds, but flasks, beakers, etc., are 
made free hand without any moulds, the glass blower taking 
advantage of the fact that the surface tension of the glass will 
cause it to contract when heated, while he can make it expand 
at will by blowing into the vessel when it is hot. Glass tubing 
is made by rapidly drawing out a hollow mass of glass. This 
is done by two men who walk or run in opposite directions. 
Pressed glass articles are made by pressing the hot glass in 
moulds having the desired designs cut upon them. The designs 
in cut glass are worked into the glass with rapidly revolving 
wheels, armed with wet sand which quickly cuts away the glass. 
The surface is afterward polished. Plate glass is made by casting 
the glass into large sheets which are ground until the sides 
are plane and parallel, and then polished. 

The cheaper kinds of soda glass are often made from sodium 
sulfate, carbon, calcium carbonate, and sand, the equation being, 

2Na 2 SO 4 + C + 2CaCO 3 + 12Si0 2 = 2Na,CaSi 6 Oi 4 '+ 3CO 2 + 

2S0 2 

Soda glasses are rather easily attacked by water, and for many 
chemical purposes the more difficultly soluble and less fusible 
potassium glass is made; this is practically the same as the 
soda glass except the potassium has taken the place of the so- 
dium. Its composition may be represented by K 2 CaSi 6 Oi 4 . It 
is commonly called hard or Bohemian glass. Curiously enough 
a glass composed of a mixture of soda and potash glass is more 
soluble than one containing either alone. 

Many special glasses are made in which the sodium, potas- 
sium, and calcium are replaced more or less completely by lead, 
magnesium, zinc, barium, antimony, arsenic, aluminum, lithium, 
didymium, thallium, iron, manganese; and the silica by boric 
or phosphoric oxides. Flint glass, crystal glass, and Strass are 
potassium lead glasses and have great brilliancy and a high index 
of refraction. They are easily attacked by water and acids, 
and are not suitable for chemical purposes, but are used in 
optical instruments. One kind of Jena glass is especially resist- 
ive to chemical action and is a borosilicate glass containing zinc 
and barium. 

Molten glass will dissolve many metallic oxides which often 


impart characteristic colors. Cobalt gives blue, chromium and 
copper oxides green, uranium yellow with a greenish fluorescence, 
cuprous oxide with a reducing agent, selenium and metallic gold 
a deep red color due to a colloidal solution of the elements. Milk 
glass is made by adding bone ash (calcium phosphate) or stannic 
oxide, Sn(>2, to the glass. These do not dissolve, but make the 
glass white and opaque. 

Glass is an amorphous substance, and h&s no definite melting 
point; it simply gradually softens as the temperature is raised. 
This has led many to consider glass at ordinary temperatures as a 
very viscous liquid, but it certainly is highly elastic and has the 
general properties of a solid. When glass is kept for a long time 
at temperatures near its softening point, it gradually becomes 
crystalline or devitrifies as it is called. With certain kinds of 
glass, this seriously interferes with its working. 

All thick glass objects, especially if of irregular shape, must be 
cooled very slowly, annealed as it is called. If cooled quickly, 
the outside will harden while the inside is still soft; upon further 
cooling, the interior tends to contract and this puts great stress 
upon the glass and makes it extremely likely to crack. A piece 
in this condition will often fly into minute fragments when 
scratched at any point on the surface. 

Analytical Properties of Calcium. The more important 
analytical properties of calcium have been given in connection 
with the oxalate. In addition it may be mentioned that volatile 
calcium compounds, the chloride for example, color the Bunsen 
flame brick red and give a red and a green band in the spectro- 
scope (see Frontispiece). The latter test is quite sensitive. 


Strontium is the second member of the alkaline earth metals, 
if they are arranged in the order of their atomic weights. 

Both the physical and chemical properties of the element and 
of its compounds closely resemble those of calcium and the 
calcium salts. It is, however, much less abundant and is rela- 
tively unimportant. 

The metal may be prepared by electrolysis of the fused 
chloride; it is silvery white and crystalline. It melts at 800 

GROUP II '331 

and volatilizes at 950 in a vacuum. The density is 2.5. It is 
more rapidly oxidized in air than calcium and decomposes 
water and alcohol. It dissolves very rapidly in acids. When 
heated in oxygen it burns with a bright red flame. It combines 
with most of the non-metallic elements. The atomic weight 
of strontium is 87.63. 

Occurrence. Strontium occurs in nature as strontianite, the 
carbonate, SrCOs, isomorphous with aragonite and as celestite, 
SrS0 4 , isomorphous with anhydrite, CaS0 4 . 

The Compounds of Strontium. Besides the naturally occur- 
ring carbonate and sulf ate, the principal compounds of strontium 
are the chloride, SrCl 2 -6H 2 0, the nitrate Sr(N0 3 )2, the oxide, 
SrO, the hydroxide Sr(OH) 2 , and the oxalate SrC 2 4 : The 
salts may be easily made from the carbonate by treating the 
latter with the acid corresponding to the salt desired. The 
oxide may be made by a similar method to that used for calcium 
oxide, that is by heating the carbonate; but the temperature 
required is far higher than that for calcium oxide, so a better 
way is to heat strontium nitrate. The hydroxide may be easily 
made by passing superheated steam over the carbonate, 

SrC0 3 + H 2 O <= Sr(OH) 2 + C0 2 

The steam sweeps away the C0 2 as it is formed and so favors 
the transformation of the carbonate. Strontium chloride, 
sulfate and carbonate are less, while the hydroxide and oxalate 
are more soluble than the corresponding calcium compounds. 

Volatile strontium compounds color the flame carmine red and 
give a spectrum which consists of a number of strong red lines 
and a weaker but more characteristic blue line. The spectro- 
scopic test is very sensitive. 


Barium is the third element of the alkaline earth family. The 
metal is silver white and has a density of 3.78. It melts at about 
850 and boils at 1,150. It oxidizes rapidly, sometimes catching 
fire spontaneously, and decomposes water and alcohol even more 
vigorously than strontium. It combines readily with hydrogen 
and nitrogen. The atomic weight of barium is 137.37. 


Occurrence. Barium occurs in nature principally as barytes 
or heavy spar which is barium sulfate, BaS0 4 , and as the car- 
bonate or witherite, BaCOs. From these the other barium com- 
pounds are made. They are technically much more important 
than the corresponding strontium compounds. Barium salts 
may be obtained by treating witherite with the corresponding 
acid or by reducing the sulfate to sulfide with carbon at high 
temperatures, and treating the sulfide with the acid of the desired 

Barium Oxide and Hydroxide. Barium oxide, BaO, is made 
by heating the nitrate to a high temperature. It is impracticable 
to prepare the oxide by heating the carbonate alone, since the 
temperature of decomposition is very high, but a mixture of 
carbon and barium carbonate reacts easily at a moderate tem- 
perature to form barium oxide and carbon monoxide, 

BaCO 3 + C = BaO + 2CO 

The formation of the carbon monoxide aids by reducing the 
partial pressure of the carbon dioxide. The oxide combines 
even more readily with water than calcium oxide to form the 
hydroxide which is often called baryta, Ba(OH) 2 . Because of 
this, it is feasible to prepare the hydroxide by heating the 
carbonate in a stream of superheated steam, 

BaC0 3 + H 2 O = Ba(OH) 2 + C0 2 

The steam decreases the partial pressure of the carbon dioxide, 
and the opportunity for the formation of the very stable hydroxide 
aids in the reaction. 

The hydroxide is more soluble in water than either the corre- 
sponding strontium or calcium compound and is a somewhat 
stronger base. It crystallizes from its solutions with 8 molecules 
of water, Ba(OH) 2 -8H 2 O. 

Barium Dioxide. Barium dioxide, Ba0 2 , is formed when 
either the oxide or hydroxide is heated in a stream of air. At a 
higher temperature, or lower pressure of oxygen, it decomposes 
into oxygen and barium oxide. This was the basis of a process 
for obtaining oxygen from the air. This process has now been 
replaced by that founded on the liquefaction of the air. 


Barium dioxide is important as the starting point for the 
preparation of hydrogen dioxide. 

Barium Carbonate. Barium carbonate, BaC0 3 , as has been 
mentioned, occurs in nature. It is also formed by bringing 
together a solution of a carbonate and a barium salt, or by 
fusing the sulfate with sodium carbonate. It is only slightly 
soluble, but more so than the carbonates of calcium or strontium. 
Its solubility in water is increased by the presence of carbon 
dioxide, owing to the formation of the bicarbonate. 

Barium Sulfate. Barium sulfate, barytes, BaS04, is in some 
respects the most important salt of barium. It is the chief 
source of the other compounds, and is itself extensively used as a 
white pigment, called " permanent white," which is used in 
paints. The precipitated barium sulfate is more valuable as a 
pigment than the natural because the particles are smaller and 
hence it has more " covering" power since it is not so transparent. 
This precipitated sulfate may be obtained by first dissolving 
witherite in hydrochloric acid, and then adding sulfuric acid 
which will precipitate the barium sulfate and regenerate the 
hydrochloric acid which is then ready to treat another lot of 
witherite. It may also be obtained by reducing the sulfate to 
sulfide, BaS, by carbon at high temperatures, and treating the 
extract from this, which will contain the hydroxide and the 
hydrosulfide, with sulfuric acid. The sulfate has some advan- 
tages over white lead as a pigment, in that it is chemically inert 
and cannot change in color, but it does not "cover" as well 
(see p. 422). 

Barium sulfate is one of the least soluble salts; 100 grm. of 
water will dissolve 0.00026 grm. of the salt at 25. 

Barium Sulfide. Barium sulfide, BaS, to which frequent 
reference has been made, is prepared by the interaction of carbon 
and the sulfate 

BaSO 4 + 4C = BaS + 4CO 

It is slowly attacked by water, forming the hydroxide and 

Barium sulfide is phosphorescent when it contains certain 
impurities, although the pure compound is not. 

Barium Chloride. Barium chloride, BaCl 2 .2H 2 O, is made by 


heating a mixture of the sulfate with carbon, and calcium 
chloride, CaCU, 

BaS0 4 + CaCl 2 + 4C = BaCl 2 + CaS + 4CO 

The furnace product is treated with water to dissolve the barium 
chloride, which is purified by recrystallization. It is not as 
soluble as calcium chloride and is not deliquescent. It is much 
used as a reagent for the detection and determination of the 
sulfate ion. Like all the other soluble barium salts it is highly 
poisonous. * 

Barium Chlorate. Barium chlorate, Ba(ClO 3 )2, is made by 
heating together, in solution, barium chloride and sodium 

BaCl 2 + 2NaC10 3 = Ba(C10 3 )2 + 2NaCl 

It is used in making green fire. 

Barium Nitrate. Barium nitrate, Ba(NO 3 )2, is made by the 
action of nitric acid upon the carbonate, sulfide, oxide or hy- 
droxide of barium. It crystallizes as the anhydrous salt and 
sometimes is used in colored fire. It is used in the preparation 
of the oxide as noted above. 

Barium Chromate. Barium chromate, BaOCX, is precipi- 
tated by bringing together a barium salt in solution and a 
soluble, chromate or dichromate. It has a yellow color and is 
one of the few colored barium salts. It is but slightly soluble 
in water and weak acids, but dissolves in strong acids. The 
chromates of calcium and strontium are very much more soluble 
than barium chromate, and advantage is taken of this in the 
analytical separation of the elements. 

Analytical Properties of Barium. The properties of barium 
which are used in analysis are the slight solubility of the carbon- 
ate, sulfate, and chromate together with the green color which 
it will impart to the Bunsen flame, and the spectrum which con- 
sists of a number of orange and green lines. These are much less 
intense than the lines of calcium and strontium so that this test 
is not very sensitive. In the separation of the group the mem- 
bers are precipitated together as the carbonates. The precipitate 
is then dissolved in acetic acid and potassium dichromate added 
which precipitates barium chromate. The solution is filtered 


and a dilute solution, 5 grm. per liter, of potassium sulfate added. 
Since strontium sulfate is much less soluble than calcium sulfate, 
the concentration of the sulfate as ion in this dilute solution will 
precipitate strontium sulfate, but not that of calcium. The solu- 
tion is then filtered and ammonium oxalate added to precipitate 
the calcium. 


The remaining member of the group is the very rare element, 
radium. This was discovered in 1898 by M. and Mme. Curie. 
It is widely distributed in nature, but always in very small 
quantities. All rocks and soils seem to contain a very little. 
It is most abundant in the minerals which contain uranium 
and there seems to be a more or less definite ratio between the 
radium and the uranium content of the ores. 

The compounds of radium very closely resemble those of 
barium, and the corresponding compounds of the two elements 
are often isomorphous. Radium chloride, bromide, carbonate, 
and sulfate are less soluble than the corresponding barium 
salts. This decrease in solubility is in accord with its higher 
atomic weight, 226.0. The volatile compounds color the 
Bunsen flame carmine red and give two beautiful bands in the 
red, a line in the blue-green and two weak lines in the violet 
portion of the spectrum. The spectrum of radium is so distinct 
and characteristic that there can be no doubt but that it is an 
element. The radium compounds have the very remarkable 
property of being highly radioactive and of changing sponta- 
neously into other elements, but the discussion of this phase of 
the subject will be postponed to a later point when the radio- 
active elements may be considered in a group. 

Metallic radium has been obtained by first preparing the 
amalgam through the electrolysis of a solution of radium chloride 
using mercury as the cathode, and then carefully driving off the 
mercury by heating in a vacuum. The radium is left behind as 
a hard silvery white metal. It melts at 700 and is rather more 
volatile than barium. It blackens in the air, due to the forma- 
tion of a nitride, acts upon water very vigorously. The radio- 
activity of the metal is the same as that of an equivalent quantity 
of one of its salts. 


General Relationships. As with the alkali metals, the proper- 
ties of the alkaline earth elements vary in a regular way with the 
atomic weight as may be seen from the following table. 

As the atomic weight increases. 

The reactivity of the element increases. 

The density of the element and compounds increases. 

The basic properties of the hydroxide increase. 

The solubility of the hydroxide increases. 

The solubility of the halogen compounds and nitrate decreases. 

The solubility of the sulfate and chromate decreases. 


The members of this sub-group, beryllium, magnesium, zinc, 
cadmium, and mercury do not form as well marked a family 
as the members of the alkaline earth metals. 

Beryllium is perhaps even more like the next group than it is 
like the other members of this and so makes a good connecting 
link between the two groups. Magnesium, zinc, and cadmium 
have many points in common but not as many as calcium, 
strontium and barium. Mercury differs markedly from the 
others as will be seen when its compounds are discussed. All 
the members of the sub-group are more readily reduced to the 
metallic state than are the members of the alkaline earths and 
are less active chemically. 


The first member of the magnesium group is beryllium or 
glucinum as it is often called. It got the name glucinum from 
the fact that its salts have a sweet taste and the name beryllium 
from its occurrence in the gem stone beryl. 

The element does not occur free in nature and its compounds 
are rare. Neither the element nor its compounds have found any 
application except that some of the latter are used as gems. 
Beryl, for example, is a double beryllium aluminum silicate, 
Be 3 Al 2 Si6Oi 8 , which when transparent and green colored is 
known as emerald and when having a bluish-green tint as 
aquamarine. The color of emerald is probably due to chromium. 

Metallic beryllium may be prepared by the electrolysis of the 
fused double fluorides of sodium or potassium, BeF^NaF. 
It is less active chemically than the metals of the alkaline earth 
group and decomposes water only very slowly, even when heated. 
It dissolves readily in the dilute acids, and in solutions of sodium 
or potassium hydroxides. Aluminum, the first metallic member 
22 337 


of the next group, shows this same peculiarity and this is but 
one of several indications that beryllium gives of its position as 
the connecting element between Groups II and III. It will be 
recalled that in discussing lithium it was pointed out that the 
first member of a group often had properties which connected 
that group with the next higher one. 

The chloride, BeCla, and the sulfate, BeSO 4 , are the best known 
salts. They are both soluble in water and their solutions react 
acid, owing to hydrolysis. The hydroxide, Be(OH)i, is very 
slightly soluble in water but dissolves in acids, and also in solu- 
tions of sodium or potassium hydroxide. 

Beryllium carbonate is very unstable and decomposes in the 
air at temperatures below the boiling-point of water into the 
oxide and carbon dioxide. 

The atomic weight of beryllium is 9.1. 


Magnesium occurs very abundantly in nature as a component 
of a large number of silicates of which mention may be made of 
olivine, MgaSiO 4 , serpentine, (MgFe) s SiiOr2H 2 O, asbestos, 
CaMg(SiOs)4, meerschaum, H a Mg 2 (SiO,) 5 H 2 O; as the car- 
bonate, magnesite, MgCO 5 , and as a double salt with calcium 
carbonate, MgCO 5 -CaCO 3 , known as dolomite. The sulfates 
and chloride and mixtures of the double salts are found in 
Stessf urt. This by no means exhausts the list of its occurrence 
which might be greatly extended. 

Preparation of the MetaL Metallic magnesium is now 
prepared by the electrolysis of fused dehydrated carnallite, 
MgCla-KCHHaO. The electrolysis is carried out in closed elec- 
tric furnaces through which a current of hydrogen or coal gas is 
passed to protect the metal from the air. 

Magnesium is a silver white metal having a density of 1.75. 
It melts at 650 and boils at 1,120. When heated nearly to its 
melting-point it softens and may be pressed into wire which can 
then be rolled into ribbon. It preserves its luster in dry air 
but tarnishes in moist. When heated in the air it burns with 
a very intense light, forming a mixture of the oxide, MgO, and 
the nitride, Mg*N* At a moderately high temperature, the 


metal decomposes steam, forming hydrogen and the oxide. It 
will burn in carbon dioxide with the deposition of carbon and 
the formation of the oxide. As may be judged from these 
examples, mangesium is an excellent reducing agent. 

The light produced when magnesium burns is especially rich 
in the violet rays which most affect the photographic plates and 
hence the metal finds application as the active component of 
"flash powder," which may be prepared from powdered mag- 
nesium and potassium chlorate or parnilar oxidizing agents. 
With aluminum, it forms useful alloys, which can easily be cast 
and worked and are lighter thn aluminum. 

The atomic weight of magnesium is 24.32. 

Magnesium Oxide and Hydroxide. ^lagnesium oxide, MgO, 
or "calcined magnesia" is prepared by heating the carbonate, 
which breaks down into the oxide and carbon dioxide at a much 
lower temperature than calcium carbonate. It is a white, highly 
infusible substance, which is used as a lining for electric furnaces 
and for crucibles for use in the Goldschmidt process, p. 396. 
It slowly reacts with water to form the hydroxide, Mg(OH), 
which is only slightly soluble. The solution has a very faint 
alkaline reaction, due to the fact that the hydroxide is a weaker 
base than calcium, as well as being less soluble. Because of its 
slight solubility the hydroxide is precipitated when a soluble 
hydroxide is added to a magnesium salt. An exception to this 
is found in the fact that if ammonium salts are present in suffi- 
cient concentration, magnesium hydroxide is not precipitated 
by ammonium hydroxide. The explanation is as follows: 
Ammonium hydroxide is a weak base and the reaction represented 
by the equation, 


Ci Ci C 5 

takes place only to a slight extent with the formation of ammo- 
nium and hydroxyl ions. From the law of mass action, CaCj/Ci = K. 
Ammonium salts are largely dissociated and give a high concen- 
tration of the ammonium ion. So, if some ammonium salt, say 
the chloride, be added- to an ammonium hydroxide solution, the 
concentration of the ammonium ion, c*, will be increased and 
this wul cause a decrease in the concentration of the hydroxyl 


ion, cs, and an increase in that of the undissociated hydroxide 
d. Ammonium hydroxide is then a weaker base in the presence 
of ammonium salts than when alone; and when the concentra- 
tion of the salt is high, that of the hydroxyl ion becomes very 
small indeed. 

When ammonium hydroxide is added to a magnesium salt, the 
reaction is represented by the following equation, 

Mg++ + 2C1- + 2NH 4 + + 20H-4=Mg(OH) 2 + 2NH 4 + + 2C1- 

c' 8 

The conditions for equilibrium are represented by the equation 
c's/c'icV = K. Now, before magnesium hydroxide can be pre- 
cipitated, the solution must be saturated with the substance; 
which means the c' 3 must have reached a certain definite value; 
and hence c'i cV must have attained a fixed and definite mag- 
nitude known as the solubility product or " precipitation value." 
When ammonium hydroxide is added to a magnesium salt in the 
absence of an ammonium salt, c'z is large enough so that c'icV 
exceeds the solubility product of magnesium hydroxide and this 
compound is precipitated; but if an ammonium salt is present, 
c'2 will be so reduced as shown above that c'icV will be smaller 
than the solubility product, and Mg(OH) 2 will not be precipi- 
tated. Correspondingly, magnesium hydroxide is dissolved by 
ammonium salts, such as the chloride or nitrate. 
This may be shown as follows: 

Mg-H- + 


^Mg(OH) 2 


c' 2 

c' 3 

2C1- + 

2NH 4 + 

<=NH 4 C1 

C ' 3 

'. n' 2 

The ammonium ion from the salt unites with the hydroxyl ions 
from the magnesium hydroxide, and thereby decreases c' 2 ; this 
will, of course, decrease c's, and hence increase the solubility of 
the hydroxide. 

Acids would, of course, act in 'much the same way and decrease 
the concentration of the hydroxyl ion through the formation of 
water. Their action, however, is more vigorous than that of 
ammonium salts. Many other hydroxides are affected in the 


same way, so that the general rule may be formulated, that 
moderately soluble hydroxides will be dissolved by ammonium 
salts and even very difficultly soluble ones by acids, in each 
case, because of the decrease in the concentration of the 
hydroxyl ion. 

Magnesium Carbonate. The normal carbonate occurs in 
nature as magnesite, MgCOa. It also combines with calcium 
carbonate to form a compound called dolomite, MgCOs, CaCOs, 
which occurs in enormous quantities. 

Magnesium carbonate is but slightly soluble in water, although 
more soluble than the hydroxide. It is more soluble in water 
containing carbon dioxide than in pure water, because of the 
formation of the bicarbonate. 

When a soluble carbonate, such as sodium carbonate, is added 
to a magnesium salt, a mixture of magnesium carbonate and 
hydroxide is precipitated. The cause of the precipitation of 
the hydroxide is found in the fact that the sodium carbonate is 
hydrolyzed, giving a certain amount of the hydrocarbonate ion, 
and of the hydroxyl ion, in addition to the carbonate. Then, 
since magnesium hydroxide is even less soluble than the carbon- 
ate, the two salts are precipitated together. 

This mixture prepared by precipitation is washed and dried 
and put on the market under the name of magnesia alba. It is 
used in medicine as a mild alkali. 

Magnesium Chloride. Magnesium chloride, MgCl 2 -6H 2 0, is 
found in the Stassfurt salts and is called bischofite; it is very 
soluble and is highly deliquescent. When it is attempted to 
dehydrate the salt by heating, decomposition takes place and 
magnesium oxide and hydrochloric acid are formed as shown in 
the following equation: 

MgCl 2 + H 2 O = MgO + 2HC1 

Magnesium chloride forms a double salt with potassium 
chloride, MgCl 2 -KCl-6H 2 0, which is known as carnallite. 

When carnallite is heated it loses its water of crystallization 
readily but the magnesium chloride is not decomposed. An 
ammonium magnesium double chloride, MgCl 2 -NH 4 Cl-6H 2 O, 
isomorphous with carnallite is known. This may be dehydrated 
without decomposition of the magnesium chloride, and when 


heated to a still higher temperature the ammonium chloride is 
driven off and anhydrous magnesium chloride is formed. 

The term double salt has been used a number of times without 
definition. By it is meant a salt formed by the chemical union 
of two salts and characterized by the fact that its solution shows 
the properties of all the ions of the component salts. 

A complex salt is similar to double salts in that it is formed by 
the chemical union of two or more salts, but differs in that its 
solutions do not exhibit the properties of the ions of the salts 
from which it was formed but instead some of these properties 
are found to have disappeared and to have been replaced by 
entirely new properties. 

A typical complex salt is potassium ferrocyanide, K 4 Fe(NC) 6 . 
This may be formed by bringing together potassium cyanide, 
KNC, and ferrous cyanide, Fe(NC) 2 , 

4KNC + Fe(NC) 2 = K 4 Fe(NC) 6 

It might be expected to show the properties of potassium, 
ferrous, and cyanogen ions, but it shows the properties of potas- 
sium as ion and of a new very stable complex ion called ferro- 

cyanogen, Fe(NC)e , and not the properties of either 

ferrous or cyanogen ions. A salt of such a complex ion is called 
a complex salt. Complex salts and double salts gradually merge 
into one another, so that the difference is one of degree rather 
than kind. 

Magnesium Sulfate. Magnesium sulfate, MgSO 4 , occurs in 
nature as kieserite, MgS0 4 -H 2 O, epsomite, MgS0 4 -7H 2 0, 
schoenite, MgS0 4 -K 2 S0 4 -6H 2 O, and in a number of other double 
salts in the Stassfurt deposits and in other places. 

The hepta-hydrate or Epsom salts is an important compound, 
being used in the manufacture of potassium sulfate from the 
chloride, as a dressing for cotton goods, and in medicine as a 
purgative. It is freely soluble in water/ and in this way differs 
strikingly from the sulfates of the alkaline earths and is like the 
sulfate of beryllium. 

Schoenite, MgSO 4 -K 2 S0 4 -6H 2 O, belongs to a rather large group 
of isomorphous salts which have the general formula, MAO 4 - 
m 2 A0 4 -6H 2 O, in which M stands for Mg, Zn, Cd, Ni, Fe, Mn, and 
m for K, Kb, Cs, NH 4 , and A for S, Se, or Cr, These salts are 


characterized by the fact that they readily lose their water of 
crystallization, in general, without the decomposition of the salt 
in other ways. 

Magnesium Sulfide. Magnesium sulfide, MgS, is made by 
the direct union of the elements. It is even more readily hydro- 
lyzed than calcium sulfide, owing to the weakness of the hydrox- 
ide and its smaller solubility. By boiling the solution, all the 
sulfur may be obtained as hydrogen sulnde. 

Phosphates of Magnesium. Magnesium ammonium phos- 
phate, MgNH 4 PO 4 6H 2 0, is the most important phosphate of 
magnesium. Because of its slight solubility it is formed whenever 
magnesium, ammonium, and phosphate ions are brought together. 
The solution must be alkaline in order to have a sufficient amount 
of the phosphate ion P0 4 present (see sodium phosphate, p. 
290), and hence the solution must contain ammonium salts to 
prevent the precipitation of Mg(OH) 2 . The precipitate is of 
course soluble in acids, since phosphoric acid is weak. This 
compound is analytically the most important of the magnesium 
salts, since magnesium is always detected and determined 
through its formation. When magnesium ammonium phos- 
phate is heated, it decomposes into magnesium pyrophos- 
phate, Mg 2 P207, ammonia and water, 

2MgNH 4 P0 4 = Mg 2 P 2 7 + 2NH 3 + H 2 O 

In quantitative analysis, magnesium is weighed as the pyro- 

Magnesium Nitride. Magnesium nitride, Mg 3 N2, is formed 
by the interaction of magnesium and nitrogen at a fairly high 
temperature. It reacts with water to form ammonia and mag- 
nesium hydroxide, 

Mg 3 N 2 + 6H 2 = 3Mg(OH) 2 + 2NH 3 

Analytical Reactions of Magnesium. Compounds of mag- 
nesium do not color the Bunsen flame, and consequently no 
spectrum is to be obtained in this way; but by passing electric 
sparks between pieces of the metal, a characteristic spectrum 
may be secured. The analytical importance of the magnesium 
ammonium phosphate has been discussed. The ion and all 


of its compounds which do not contain colored anions are 

Water Softening. Natural waters almost invariably con- 
tain bicarbonates, chlorides, sulfates, and occasionally nitrates 
of calcium, magnesium, sodium, potassium, iron, and aluminum 
and also some free carbonic acid and silica. Of course, they 
sometimes contain other salts and acids in addition. For 
example, free sulfuric acid is often found in mine waters and 
in river water in the vicinity of factories using this acid. 

Waters which contain calcium and magnesium salts are called 
hard waters and are a source of considerable annoyance and 
expense in the household, owing to the fact that calcium and 
magnesium soaps are practically insoluble, and that before a 
lather can be produced, soap must be added equivalent to all 
such salts which are present. The calcium and magnesium soaps 
so produced form a dirty disagreeable scum and are objectionable 
from every point of view. Many of the calcium and magnesium 
salts also cause a great deal of trouble and expense through the 
formation of " scale" in steam boilers, which not only increases 
the cost of repairs, but also causes a great waste of fuel. Magne- 
sium salts are somewhat hydrolyzed, owing to the weakness of 
magnesium hydroxide and hence are acid in reaction. This is 
often the cause of serious pitting of boilers. 

It is very desirable, therefore, that the salts of calcium and 
magnesium be removed from the waters before they are used 
for domestic or boiler purposes. This may be readily done at a 
cost far less than that of the soap, repairs and extra fuel, which 
would otherwise be required. For the removal of the objection- 
able salts, advantage is taken of the fact that calcium car- 
bonate and magnesium hydroxide are nearly insoluble in water, 
and the treatment is so managed that these salts shall be pre- 
cipitated. This is done by adding the proper amount of slaked 
lime and of crude sodium carbonate, or soda ash as it is called, 
to the water. The quantities to be added are calculated from 
the analysis of the mineral content of the water. Before treat- 
ment, the average water will contain free carbonic acid, bicar- 
bonates, sulfates, and chlorides of magnesium, calcium, and 
sodium or more correctly, it will contain the corresponding ions. 
After treatment it will contain sulfate and chloride of sodium 


and the slight amount of calcium carbonate and magnesium 
hydroxide corresponding to the solubilities of these compounds, 
which is so small that it may be neglected. 

Enough sodium carbonate is added so that the total of the 
sodium ion originally present plus that added shall equal in 
equivalents the sulfate and chlorine ion present in the untreated 
water. Calcium hydroxide is added in sufficient amount to 
neutralize any free acid, to change all the hydrocarbonate ions 
to carbonate and to precipitate the magnesium as hydroxide. 

What promises to become an important method for softening 
water is based upon the fact that when a hard water is filtered 
slowly through a bed of artificial zeolite, NaAlSi0 4 -3H 2 O, called 
"permutite" the calcium, magnesium, manganese, and iron are 
automatically removed and replaced by sodium from the zeolite. 
The filtering medium becomes exhausted, but may be rejuve- 
nated by treatment with a strong solution of sodium chloride 
which reverses the action and dissolves the calcium, magnesium, 
etc., replacing them by sodium. 


General. Starting with beryllium, zinc is the third member 
of the magnesium sub-group. It does not occur as abundantly 
in nature as magnesium, but is of far more importance tech- 
nically. The world's annual production of metallic zinc is 
about 880,000 tons, of which the United States furnishes about 
30 per cent. Kansas, Missouri, Wisconsin and New Jersey are 
the chief zinc producing states. 

Occurrence. The principal ores of zinc are the sulfide, ZnS, 
known as sphalerite, blende, rosin-jack, or black-jack; the car- 
bonate, ZnC0 3 , called smithsonite, zinc spar and sometimes, 
though improperly, calamine; the silicates, calamine, Zn 2 H 2 Si05 
and willimite, Zn 2 SiO 4 ; and franklinite, (FeZnMn)(FeO 2 ) 2 . 

Metallurgy. The ores are first concentrated by crushing and 
washing away the lighter rock materials with water or in the case 
of the sulfide, by the recently devised flotation process which is 
rapidly increasing in popularity. In the latter process the ore 
is finely crushed and thoroughly churned up with air, oil and 
water. The ore particles, oil globules, and bubbles of air unite 



to form an aggregate which is lighter than water and so floats 
off as froth while the worthless rock sinks to the bottom. The 
ores which contain the silicates or oxides of zinc merely require 
drying to remove the water before they are ready for reduction, 
but the carbonate should and the sulfide must be converted into 
the oxide before further treatment. The carbonate changes 
to the oxide when heated to a very moderate temperature, being 
like magnesium carbonate in this respect. The sulfide must be 
roasted in the presence of air when the oxide and sulfur dioxide 
are formed: 

2ZnS + 30 2 = 2ZnO + 2SO 2 

A part of the sulfur dioxide so produced is used to make sulfuric 
acid, and now furnishes about 250,000 tons of this substance per 



After getting the zinc in the form of anhydrous silicate or 
oxide, the next step is to reduce with carbon in a fire-clay retort 
furnace, Fig. 58. The equation for the chief reaction is 

ZnO + C = Zn + CO 

The temperature of reduction is about 1200, while the boiling- 
point of zinc is 906; hence, the zinc will be vaporized as rapidly 
as formed, and provision must be made to catch and condense the 

When the condenser is cold, that is below 419, the melting- 


point of zinc, the metal collects in the form of a fine powder which 
contains a few per cent, of the oxide, and is known as zinc dust; 
when it is above 419, the zinc collects in the liquid state and is 
tapped off and cast into moulds forming what is known as spelter. 

Under the best working conditions, the loss is great, as indi- 
cated by the bluish-green color of the flames burning at the 
mouths of the condensers, and the great amount of zinc oxide 
fumes which are given off and lost. An electrolytic process has 
recently been developed to the commercial stage in this country. 
The ore is roasted at a low temperature which produces ZnO 
and ZnSO4. The product is extracted with dilute sulfuric acid, 
and the resulting impure solution of zinc sulfate is purified and 
then electrolyzed, the zinc being deposited at the cathode. It 
is more than 99.99 per cent, pure and commands a premium on 
account of its purity. 

Physical Properties. A fresh surface of pure zinc is brilliantly 
white, but quickly tarnishes and then has the familiar bluish- 
gray tinge. It melts at 419 and boils at 906. Cast zinc is 
crystalline and very brittle, but when heated to between 100 
and 150, it becomes malleable and may be rolled into thin sheets 
or drawn into wire. At about 200 it again becomes brittle and 
may be powdered in a mortar. The density of the cast zinc is 
about 6.93 and of the rolled 7.18. 

Chemical Properties. Zinc belongs to the more active of the 
elements, but is far less so than the alkali or alkaline earth 
metals. It is not affected by dry air at ordinary temperatures, 
but burns brilliantly above 500. In ordinary air which contains 
water vapor and carbon dioxide, it quickly becomes covered 
with a basic carbonate which acts as a protective coating and 
keeps the rest of the metal from further corrosion. 

The behavior of zinc toward dilute sulfuric acid is very peculiar 
in that while the commercial metal dissolves easily, very pure 
zinc is not attacked until it is touched by a piece of copper, 
platinum, carbon, etc., then the evolution of hydrogen begins, 
but on the copper, etc., and not on the zinc. However, the zinc 
passes into solution while the copper is unchanged. The explana- 
tion is found in the fact that it is very difficult for hydrogen as 
ion to change to hydrogen gas at the surface of zinc; so much so 
that the action is prevented with the pure metal. But if copper 


is in contact with the zinc, the change of hydrogen as ion to hy- 
drogen gas takes place easily on the copper, and the positive 
charges of the hydrogen ion are conducted from the copper to 
the zinc and change the latter into the zinc ion. 

Sodium or potassium hydroxide will slowly dissolve zinc 
forming hydrogen and zihcate: 

2NaOH + Zn = Na 2 ZnO 2 + H 2 

Zinc is used in making brass, which is 63 to 70 per cent, 
copper and 37 to 30 per cent, zinc; in many other alloys; in 
electric batteries; and especially in galvanizing iron. 

Galvanized Iron. Two-thirds of the zinc produced per year 
goes into galvanized iron, that is, iron coated with a layer of 
zinc. This resists the action of the weather better than plain 
iron, because the zinc oxidizes first and so protects the iron. 
The iron is coated with zinc in three ways : first, by dipping very 
clean sheets of iron into molten zinc; second, by electrically 
depositing zinc upon cathodes of iron from properly prepared 
solutions of zinc salts; third, by the process known as sherard- 
izing, which consists in exposing the cleaned iron object to the 
vapors of metallic zinc. In some ways this last method has the 
advantage, because the zinc and iron adhere better than they 
do by the other methods. 

Zinc Oxide and Hydroxide. Zinc oxide, ZnO, or zinc white, 
is by far the most important manufactured compound of zinc, 
and is prepared on a very large scale, 59,000 tons p'er year, 
either directly from the ores or from the metal. In this country 
the ores are used exclusively and the process is briefly as follows : 
the ore, either franklinite or roasted sulfide, is mixed with 
anthracite coal and charged into a furnace upon a brisk fire of 
hard coal with a good draft of air; the zinc oxide is reduced, the 
metal volatilizes, and then is reoxidized, forming a very fine white 
powder which is carried along with the flue gases until well cooled, 
and is then collected in cloth bags which act as filters. The prod- 
uct is used as a pigment in paints, in rubber goods, as a start- 
ing point for the manufacture of other zinc compounds, and 
as a base for ointment in medicine. Its largest single use is in 
paints. It is not as poisonous as white lead, and does not blacken 
with hydrogen sulfide. A paint containing zinc oxide as the 


only pigment gets very hard and brittle, so it is unsuited for 
exterior work; but a mixture of white lead (basic lead carbonate) 
and zinc white is a great favorite and lasts better than either 
alone. Rubber goods consume a great deal of zinc oxide, 
automobile tires alone are said to require 40,000,000 Ib. a year. 

Zinc oxide is yellow when hot and white when cold. The 
change is due to a shift in the region of the absorption of light 
from the invisible to the visible part of the spectrum. 

The hydroxide is not formed by the combination of the 
oxide and water, but may be obtained by adding a soluble 
base to a zinc salt. It is a white difficultly soluble substance 
which dissolves either in acids or in an excess of sodium or 
potassium hydroxide. This is due to the fact that zinc hydroxide 
not only yields zinc and hydroxyl as ion, but* also dissociates 
to give hydrogen and zincate, Zn0 2 , ions: 

Zn+ + + 20H~ * Zn(OH) 2 <= 2H+ + ZnO a ~~ 

Ci C 2 C 3 C4 Ci 

This action is typical of a number of hydroxides and hence 
calls for a rather full explanation. We have in these cases simul- 
taneous and dependent equilibria between the hydroxide and the 
two sets of ions. The mass law equations in this case are 

c 3 . c 4 2 c 5 T ^, 
2 = K and = K' 

CiC 2 2 C 3 

If an acid is added, its hydrogen ion will combine with the 
hydroxyl and thus reduce c 2 and hence decrease c 3 making the 
solution unsaturated and causing Zn(OH) 2 to dissolve and in- 
creasing c. If on the other hand, a strong base is added the 
hydroxyl from the base will combine with the hydrogen ion from 
the zinc hydroxide and hence reduce c 4 thus causing c 3 to de- 
crease and the precipitate to dissolve to form a solution containing 
zincate ion, ZnO 2 . 

Zinc hydroxide is also soluble in excess of ammonia, due to the 
formation of a complex zinc ammonia ion, 

Zn++ + 20H- + nNH 3 = Zn(NH 3 ) n + + + 20R- 

Zinc Chloride. Zinc chloride, ZnCl 2 , may be made by the com- 
bination of the elements; by the action of hydrochloric acid 
on zinc, or on the oxide or the carbonate. It is a white, very 


deliquescent substance, which has a caustic action on the tissues. 
When a concentrated solution of the salt is mixed with zinc 
oxide, the whole sets to a hard mass of the oxy chloride, Zn(OH)Cl, 
hence the mixture is used as a cement. Solutions of the chloride 
have been successfully employed in the treatment of wood to 
prevent decay and in other cases as an antiseptic and a deodorant. 

Zinc chloride either in solution or in the fused state has the 
power of dissolving the oxides of the metals; so it is used as a 
flux in soft soldering to clean the surfaces to be joined by the 

Zinc Sulfate. Zinc sulfate, ZnSO4'7H 2 O, occurs in a hydrated 
form in nature to a small extent. It is known commercially 
as zinc vitriol or white vitriol. It may be prepared by the action 
of sulfuric acid on zinc oxide or metallic zinc or by very careful 
roasting of ZnS. It is a colorless salt, very soluble in water. It 
finds considerable use in medicine, in dyeing, in the manufacture 
of glue, and in the preparation of a zinc sulfide and barium sul- 
fate mixture known as lithophone which has been used as a 
white pigment, but is not durable. 

BaS + ZnS0 4 = BaSO 4 + ZnS 

Zinc Carbonate. Zinc carbonate, ZnCO 3 , occurs in nature 
as smithsonite, and is of considerable importance as an ore of 
zinc. A basic carbonate is precipitated upon the addition of a 
soluble carbonate to a solution of a zine salt. The term basic 
salt is applied to a mixed salt containing hydroxyl or oxygen and 
some other anion; frequently the basic salts are mixtures of the 
hydroxide or oxide with the other salt, but very often they are 
true chemical compounds. 

Zinc Sulfide. The occurrence of zinc sulfide, ZnS, in nature 
has already been discussed. It is commonly known as blende 
and in its pure state is nearly white, but it shades to brown and 
even black depending upon the impurities it contains. It may 
be precipitated by the addition of ammonium sulfide or hydrogen 
sulfide to solutions of zinc salts. In this form, it is a fine white 
amorphous powder and is used as a pigment. Of the ordinary 
heavy metals, zinc is the only one that forms a white sulfide. 
It is soluble in dilute acids, and hence must be precipitated from 
neutral solutions. If the precipitation is to be complete, the 


concentration of the hydrogen ion must be kept low. This 
may be done by adding an acetate to the solution. Most of the 
hydrogen is withdrawn in this way from the solution to form 
undissociated acetic acid. Zinc sulfide is infusible and very 
difficultly volatile. Its most important use is in the extraction 
of the metal. 

Analytical Properties of Zinc. Analytically, zinc is in a group 
with manganese, nickel and cobalt. These all have the common 
property of being precipitated as their sulfides by ammonium 
sulfide, but not by hydrogen sulfide in acid solution. Zinc and 
manganese sulfides are soluble in dilute sulfuric acid, while 
cobalt and nickel sulfide are practically unaffected. Zinc is 
distinguished from manganese by the fact that its hydroxide is 
soluble in an excess of potassium hydroxide, while that of 
manganese is not. From this alkaline solution, hydrogen 
sulfide will reprecipitate white zinc sulfide. The color of the 
sulfide is very characteristic, for it is the only common sulfide 
of a heavy metal which is white. The facts that the oxide is 
yellow when hot and white when cold, and that a green mass is 
formed when the oxide is wet with cobalt nitrate solution and 
then strongly heated are used in analytical work. 


Cadmium is an element which is very closely allied to zinc in 
its properties and is found in nature in comparatively small 
amounts associated with this metal. The boiling-point of cad- 
mium is much lower than that of zinc, consequently it is present 
quite largely in the first portions of the distillate obtained in 
the preparation of the latter metal. Jt may be freed from the 
zinc by repeated distillations at as low a temperature as possible, 
in an atmosphere of hydrogen to prevent its oxidation. This 
product may be further purified by electrolysis, a solution of 
CdSO 4 being used as the electrolyte. 

The element thus obtained is a bluish-white soft metal of a 
crystalline structure, density 8.64, melting at 321 and boiling 
at 766. The molecular weight as determined from the vapor 
density is very nearly 112.4, the number which represents the 
atomic weight, consequently the formula is Cd. 


Cadmium forms but the one simple ion, Cd + + , which is color- 
less and very poisonous to all organisms. This ion unites with 
hydroxyl to form the difficultly soluble hydroxide, a white com- 
pound, which is not redissolved on the addition of an excess of 
sodium hydroxide, but which does dissolve in ammonium 
hydroxide, due to the formation of the complex ion, Cd(NH 3 )4 + " f '. 

The oxide is a brown powder which may be formed by simply 
heating the hydroxide or by heating the metal somewhat strongly 
in the presence of oxygen. This substance dissolves readily 
in acids and may be used in the preparation of the various salts. 
Because of its color, the presence of cadmium oxide in zinc oxide 
is detrimental to the use of the latter in white paints. 

The salts of cadmium show marked resemblance to those of 
the other members of the magnesium group; the sulfate, however, 
crystallizes as shown by the formula 3CdS04'8H 2 O and is not 
analogous to ZnSO 4 -7H 2 O or MgSO 4 '7H 2 O. This salt is used 
in medicine, but finds more extensive application in the manu- 
facture of "standard cells" for electrical measurements. 

Although cadmium sulfate, nitrate, etc., are as highly dis- 
sociated as the corresponding salts of other divalent cations, the 
halogen compounds are but slightly ionized. Zinc shows a 
tendency in this same direction, and the effect is most pro- 
nounced with mercury. 

Cadmium Sulfide. Cadmium sulfide, CdS, occurs in nature as 
greenockite and is formed when hydrogen sulfide is passed into 
a solution of a cadmium salt. This produces a bright yellow 
precipitate which is slightly soluble in strong acids, but is thrown 
down quantitatively from alkaline solutions. Because of its 
fine yellow color it is used as a pigment under the name of 
cadmium yellow. It is not soluble in yellow ammonium sulfide, 
but does dissolve in the presence of a high concentration of the 
halogen ions due to the formation of the undissociated cadmium 
halide. The precipitate is more soluble in solutions of iodides 
than in chlorides, because cadmium iodide is even less dissociated 
than cadmium chloride. 

Cadmium hydroxide, also, is soluble in solutions of the alkaline 
halides for the same reason as the sulfide. 

Analytical Properties. Analytically, cadmium is in a group 
with mercury, lead, bismuth, and copper; in the separation, a 


solution is finally obtained which contains both copper and 
cadmium. These are separated by taking advantage of the 
fact that cadmium sulfide is soluble in strong solutions of com- 
mon salt while copper sulfide is not or else that copper sulfide 
is not precipitated from solutions containing potassium cyanide 
while cadmium sulfide is. 


Mercury or quicksilver as it is often called has been known 
for so long that the name of its discoverer is lost in antiquity. 

Occurrence. It occurs free in nature to a limited extent, but 
the greater part of the world's supply is obtained from the native 
sulfide, cinnabar, HgS, which is a red crystalline substance. 
The most important mines are in Spain, Italy, Austria, and the 
United States. In the United States, California and Texas are 
the principal producing states. The metal is obtained from the 
sulfide by roasting in air, when sulfur dioxide and vapors of 
mercury are formed; the latter are then condensed in air or water- 
cooled chambers. The equation for the reaction is: 

HgS + 2 = Hg + S0 2 

The great simplicity of the chemistry of this process is due to 
the instability of mercuric oxide at the high temperature of 

It is purified by treatment with 5 per cent, nitric acid, or with 
sulfuric acid and potassium dichromate, or, if it must be very 
pure by distillation under reduced pressure in the presence of a 
little oxygen to oxidize the more easily oxidizable impurities. 

Physical Properties. Mercury is a silver-white liquid metal 
which freezes at -39 and boils at 357. Mercury vapor is 
transparent and colorless and does not conduct electricity at 
temperatures near its boiling-point, but it becomes a good con- 
ductor at higher temperatures. This property is used in the 
mercury vapor lamps and rectifiers. The density of the liquid 
is 13.596 at 0. 

The density of the vapor shows that the molecular weight of 
mercury is 200, the same as its atomic weight. 

The metal is much used in physical and chemical apparatus 



and in the extraction of gold and silver from their ores, since it 
readily forms amalgams with these metals. Amalgams, as the 
alloys of mercury are called, are formed by nearly all of the 
metals. Iron and platinum are generally exceptions, but even 
these may be amalgamated *under special conditions. 

Pure mercury does not wet glass but runs around on it in clean 
round drops. When it contains even a small quantity of foreign 
metal it becomes covered with a film of oxide and then when it 
runs on glass, it strings out or forms a tail. This is a fairly sensi- 
tive test for impurities in the mercury. 

It is interesting to note that mercury is the only metal which 
is liquid at ordinary temperatures, although cesium melts at 
temperatures which are often reached during the summer. 

Chemical Properties. Mercury does not oxidize upon exposure 
to the air at ordinary temperatures, but does so very slowly at 
about 300, forming mercuric oxide. This fact was of great 
importance in the development of the present conception of 
combustion. (See p. 16.) 

Mercury combines readily with sulfur and the halogens. It 
is not attacked by dilute hydrochloric or sulfuric acids, since the 
hydrogen as ion from these is not a strong enough oxidizing 
agent to oxidize the mercury to ion. It is dissolved by nitric 
acid or by hot sulfuric acid with the reduction of a part of the 
acids, and the evolution of nitric oxide or sulfur dioxide instead 
of hydrogen. 

The compounds of mercury which are soluble enough to be 
absorbed by the human system are decidedly poisonous, but 
mercury in large drops is not. If, however, the mercury 
is very finely divided as it is in "blue mass" or in the form of 
vapor it exhibits its poisonous action. Mercury compounds 
are extremely poisonous toward bacteria and enough of the metal 
dissolves in pure water to kill these organisms in a short time. 
Mercury forms two series of compounds. In the one, it is diva- 
lent; and in the other it is at least apparently monovalent. The 
monovalent compounds are called mercurous and the divalent 
mercuric salts, and it is through the latter that the element 
establishes its connection with the magnesium group. The 
mercurous compounds will be discussed first. 



Mercurous Oxide. Mercurous oxide, Hg20, is a blackish- 
brown substance which is formed by the action of an excess of 
sodium or potassium hydroxide upon a mercurous salt. It is not 
soluble in water and is very unstable, breaking down under the 
action of light or moderately high temperature into mercury and 
mercuric oxide. Mercurous hydroxide is unknown. 

Mercurous Chloride. Mercurous chloride or calomel, HgCl or 
Hg 2 Cl2, is the most important mercurous salt. It is very diffi- 
cultly soluble in water and may be precipitated by adding a 
solution containing chlorine ions to a solution of mercurous 
nitrate. It is usually prepared by subliming a mixture of 
mercuric chloride, HgCl 2 , and mercury, 

HgCl 2 + Hg = 2HgCl 

Mercurous chloride or calomel is largely used in medicine 
because it stimulates the liver and other organs producing the 
secretions. The vapor pressure of mercurous chloride reaches 
one atmosphere below the melting-point, and therefore the 
substance sublimes. Unless very pure, it should be protected 
from light which decomposes the ordinary salt, forming mercury 
and mercuric chloride which is very poisonous. 

Mercurous Bromide. Mercurotis bromide, HgBr, is precipi- 
tated, as a white powder, by adding a bromide to mercurous 

Mercurous Iodide. Mercurous iodide, Hgl, may be formed 
directly from the elements or by precipitation from mercurous 
nitrate by an iodide. It is unstable and changes into mercury 
'and mercuric iodide. The change takes place easily in the 
presence of an excess of KI, since the latter dissolves the mercuric 
iodide formed. 

Mercurous Nitrate. Mercurous nitrate, HgN0 3 -H 2 0, is the 
most important soluble mercurous salt. It is made by acting 
upon an excess of mercury with cold dilute nitric acid. It 
hydrolyzes, forming a basic nitrate, Hg 2 (OH)NO 3 , which is but 
slightly soluble. Hence, a clear solution of the salt must contain 
nitric acid. Some metallic mercury should be kept in the solu- 
tion to reduce any mercuric ion which might form, 


Hg + Hg+ + = 2Hg+ 

Mercurous Sulfate. Mercurous sulfate, Hg 2 S0 4 , may be pre- 
pared by the action of hot-concentrated sulfuric acid upon an 
excess of mercury, or by adding the dilute acid to a mercurous 
nitrate solution. It may also be obtained by the electrolysis 
of dilute sulfuric acid using mercury as anode. It is white and 
but slightly soluble in water, but more so than the chloride. 
It is used in the construction of standard cells for the comparison 
of electromotive forces. 

Mercurous Sulfide. Mercurous sulfide, Hg 2 S, is very unstable 
and can exist only at temperatures below 10; at ordinary 
temperatures it decomposes into mercuric sulfide and mercury. 


Mercuric Oxide. Mercuric oxide or red precipitate, HgO, 
is usually prepared by heating an intimate mixture of mercuric 
nitrate and mercury until no more red fumes of nitrogen peroxide 
are given off. 

Hg(N0 3 ) 2 + Hg = 2HgO + 2N0 2 

When a solution of potassium or sodium hydroxide is added 
to a mercuric salt, a yellow precipitate of very finely divided 
mercuric oxide is formed. This precipitate is more active 
chemically and therapeutically than the coarser red oxide. 

When mercuric oxide is heated it darkens, becoming almost 
black, and at a red heat it is decomposed into the elements. It 
will be recalled, it was through this reaction that Priestly dis- 
covered oxygen. Mercuric oxide is very poisonous, but is used 
in medicine. 

Mercuric hydroxide is so unstable that it has not yet been 
prepared. The mercuric salts of the oxy-acids are highly 
hydrolyzed, so it must be a very weak base. 

Mercuric Chloride. Mercuric chloride or corrosive sublimate, 
HgCl 2 , may be made by the direct combination of the elements 
at a slightly elevated temperature. The combination takes place 
with a peculiar green flame. 

The salt is usually obtained as a crystalline sublimate by 
heating a mixture of mercuric sulfate, common salt, and a little 


manganese dioxide, the latter being added to oxidize any mer- 
curous sulfate which may be present. - 

It is much more soluble in hot water than in cold. It is soluble 
in alcohol and ether and a number of other organic solvents. 
Its aqueous solution is slightly acid, indicating some hydrolysis, 
but no basic salt is deposited as is the case with mercuric nitrate 
and sulfate. The explanation for this is furnished in the fact 
that as indicated by the electrical conductivity, the halogen 
compounds of mercury are but slightly ionized, while the sulfate 
and nitrate and the salts of other oxyacids are strongly disso- 
ciated. The inevitable results of this would be that the mercuric 
salts of the oxyacids would hydrolyze further than the halides. 
In this connection it will be recalled that the halogen compounds 
of cadmium were less dissociated than the salts of the oxyacids. 

Mercury compounds are highly poisonous to all forms of life, 
and mercuric chloride is much used as an antiseptic. For .this 
purpose it is often put up in tablets with sodium chloride which 
makes it dissolve more rapidly. It is used in medicine both 
internally and externally, but because of its highly poisonous 
nature must be used very cautiously. Its antidote is white of 
egg, with which it forms a difficultly soluble compound. This 
should be removed by a stomach pump and magnesium sulfide 
administered to precipitate the last traces of mercury. 

Mercuric Bromide and Iodide. Mercuric bromide, HgBr 2 , is 
much like the chloride, but is less soblule in water; it is also less 
dissociated than the chloride. Mercuric iodide, HgI 2 , is but 
slightly soluble in water and is formed as a yellow precipitate 
which rapidly becomes scarlet when a solution of an iodide is 
added to a solution of a mercuric salt. It dissolves freely in 
solution of potassium iodide forming a colorless very soluble 
complex salt, K 2 HgI 4 , which fails to give most of the reactions 
of mercury; for example, potassium hydroxide does not pre- 
cipitate mercuric oxide. After the addition of potassium hy- 
droxide, the solution is called Nessler reagent, and is very 
sensitive toward ammonia which gives a yellow precipitate hav- 
ing the formula Hg 2 NI-H 2 0. 

Mercuric iodide exists in the two forms; the scarlet, stable 
from ordinary temperatures to 126, and the yellow, stable 
above this temperature up to the melting-point, 253. 


Mercuric iodide and its complex potassium salt are used in 

Mercuric Nitrate. Mercuric nitrate, Hg(N03) 2 , is prepared by 
dissolving the metal in hot nitric acid, and boiling until a portion 
of the solution does not yield a precipitate of mercurous chloride 
upon the addition of a chloride. It forms colorless very deli- 
quescent crystals having the composition Hg(NO 3 ) 2 -H 2 O, easily 
soluble in water. It is easily hydrolyzed, precipitating a yellow 
basic salt. 

Mercuric Sulfate. When mercury is heated with concentrated 
sulfuric acid, mercuric sulfate, HgSC>4, water, and sulfur dioxide 
are formed. The mercuric sulfate so obtained is a white crystal- 
line salt which is easily soluble in water containing sulfuric acid, 
but hydrolyzes in pure water forming a yellow basic salt, 
Hg 3 (S0 4 )02. 

Mercuric Sulfide. The most stable compound of mercury 
is the sulfide, HgS. This exists in the two forms, black amor- 
phous and red crystalline. The latter occurs in nature as cin- 
nabar and is the most important ore of mercury. As is the gen- 
eral rule, the amorphous form is the less stable and the more 
soluble of the two, but both are exceedingly slightly soluble 
substances. The black form is precipitated whenever hydrogen 
sulfide or a soluble sulfide acts upon a mercuric salt. It is some- 
what soluble in sodium and potassium sulfide forming in the case 
of the sodium sulfide a white crystalline salt, Na 2 HgS 2 8H 2 O. 
From this solution the less soluble red crystalline form is gradually 
deposited. The red crystalline modification may also be ob- 
tained by subliming the black. The artificial red sulfide is 
called vermilion and is used as a pigment. 

Both forms of mercuric sulfide are very stable and are not 
attacked by hydrochloric acid or by dilute nitric acid hot or 
cold. The compound is dissolved by aqua regia, the sulfur ion 
being liberated as free sulfur or oxidized to the sulfate. Mer- 
curic sulfide is the most stable sulfide of the metals and this is 
utilized in the analytical separation and identification of the 

Mercuric Cyanide. Mercuric cyanide, Hg(NC) 2 , resembles 
the halogen compounds of mercury in many ways. It is more 


soluble than the halides, but is even less dissociated than these, 
and its solution is a very poor conductor of electricity. 

When heated, mercuric cyanide decomposes into mercury and 
cyanogen, C 2 N 2 , 

Hg(NC) 2 = Hg + C 2 N 2 

and this furnishes a very convenient method for the preparation 
of cyanogen. 

Mercury Fulminate. When mercury is dissolved in hot 
nitric acid and alcohol added to the still hot solution, a white 
precipitate of mercury fulminate, Hg(ONC) 2 , is soon formed. 
This, when dry, explodes upon being struck, and is used in 
percussion caps and in the caps for exploding guncotton and 

Complex Compounds of Mercury. Mercury forms a very large 
number of complex compounds far too large for them all to be 
considered in detail or even mentioned in a book of this kind. 
The nitrogen compounds, however, demand a little attention. 

Mercuric Ammonia. The mercuric salts react with ammonia 
to form compounds which may be grouped into three classes: 
first, those with ammonia of crystallization similar to salts with 
water of crystallization; second, ammono basic salts analogous 
to ordinary basic salts with NH 2 in place of OH or either NH or N 
in place of 0; third, mixed aquo ammono basic salts containing 
both OH or O and NH 2 or N. 

The first class with ammonia of crystallization may be pre- 
pared by the action of ammonia in the gaseous or liquid state 
upon mercuric salts, or by the action of aqueous ammonia upon 
these salts in the presence of ammonium salts. As examples of 
these compounds the following may be given: HgCl2'2NHs 
[fusible white precipitate], HgCl 2 -12NH 3 . 

The second class, ammono basic mercuric salts, may be formed 
by the action of aqueous ammonia upon mercuric salts in the 
absence of ammonium salts. Examples of these salts are 
HgNH 2 Cl [infusible precipitate], and Hg(NH 2 )NO 3 . 

The third class, the mixed aquo ammono basic salts, are 
formed under conditions which are more favorable to hydrolysis 
than those for the two preceding classes. The compound 
Hg 2 (OH)(NH)Cl may be cited as an example. 


Analytical Reactions of Mercury. Both the mercurous and 
mercuric ions are colorless, and the former is distinguished from 
all others by the fact that it will form a slightly soluble white 
precipitate of mercurous chloride with the chlorine ion .which 
precipitate turns black when treated with ammonium hydroxide. 

Both the mercuric and mercurous ions give mercuric sulfide 
with hydrogen sulfide in acid solution. This precipitate is 
insoluble in hot dilute nitric acid, but is dissolved by aqua regia. 
Mercuric salts are reduced by stannous chloride to mercurous 
salts or even metallic mercury as shown in the following equations, 

2HgCl 2 + SnCl 2 = 2HgCl + SnCl 4 

HgCl 2 + SnCl 2 = Hg + SnCl 4 

When metallic copper is introduced into a solution of a mercury 
salt the copper is oxidized to copper as ion and the mercury ion 
is reduced to metallic mercury which forms a bright mirror-like 
coat on the remaining copper. 

When the mercury salts are heated with lime or sodium 
carbonate, they give metallic mercury. 

General Relations of the Group. As is evident from the 
discussion given above, the members of the magnesium group 
do not form as close a family as do the members of the alkali and 
alkaline earth metals, but still there are some general relationships 
throughout the group. The increasing stability of the sulfides 
with increasing atomic weight is noticeable. Beryllium and 
magnesium sulfides are not formed in the presence of water even, 
while zinc sulfide is stable toward weak acids, but is not stable 
in contact with any but very dilute solutions of the strong acids. 
Cadmium sulfide is stable in the presence of dilute solutions 
of the strong acids, but is dissolved by concentrates! solutions 
of them and by hot dilute solutions, while mercuric sulfide is 
stable even against boiling dilute nitric acid and is scarcely 
attacked by this acid in the concentrated state. 

Beginning with zinc, the halogen compounds show a regular 
decrease in their degree of dissociation with increasing atomic 
weight both of the halogen and of the metal. The melting- and 
boiling-points of the members of the sub-group become lower as 
the atomic weights increase. 


General. Copper, silver, and gold occupy the lower part 
of the right-hand column of Group I of the periodic system. 
From their position, they should bear some resemblance to the 
alkali metals, but from what has been learned concerning the 
relations between the members of the alkaline earth metals and 
magnesium sub-group it is evident that the resemblance need not 
be close. As a matter of fact the divergence is so great that it is 
necessary to search for points of resemblance. Perhaps the 
most striking point of similarity is that copper, silver and gold 
each forms a series of compounds in which it is monovalent. In 
addition copper forms another series in which it is divalent and 
gold one in which it is trivalent. The following table will give 
at a glance a comparison of the more prominent properties of the 
two sub-groups: 

The Alkali Metals The Copper Family 

Low density 0.59-2.4. High density 8.94-19.4. 

Very active chemically. Decidedly inactive. 

Never occur free. Often found free. 

Halogen compounds are soluble. Nearly all of the halogn compounds 

of the monovalent series are but 

slightly soluble. 

Strong bases. Weak bases, except silver. 

Do not form complex ions. Form many complex ions. 


History. Copper has been known from prehistoric times and 
since it is found free in nature in abundance, it was probably the 
first metal to be put to practical use by man. Before the ex- 
tensive use of iron, tools were made of bronze, an alloy of copper 
and tin. 

Occurrence. Copper occurs in nature both free arid in com- 
bination. The native copper is sufficiently abundant so that it 




constitutes a very important ore as may be judged from the fact 
that the region in Michigan near Lake Superior produces from 
the native metal about one-sixth of the copper mined in this 
country per year. The compounds which are used as ores may 
be divided into two classes: the oxidized ores and the sulfide ores. 
The oxidized ores include cuprous oxide, Cu 2 O, which is called 
cuprite or red oxide; cupric oxide, CuO, called melaconite or 
black oxide; and the basic cupric carbonates, malachite, CuCO 3 - 
Cu(OH) 2 , and azurite, 2CuCO 3 -Cu(OH) 2 . The sulfide ores 

include cuprous sulfide, Cu 2 S, or 
chalcocite and the double sulfides of 
copper and iron known as chalcopy- 
rite, CuFeS2, and bornite, .CusFeSs. 
The sulfide ores are the most impor- 
tant because from them this country 
produces about two-thirds of its out- 
put per year. The total production 
in 1915 was nearly 1,425,000,000 Ib. 

Metallurgy. The method adopted 
for the extraction of copper from its 
ores naturally depends upon the sub- 
stances present. If the copper is 
native, all that is necessary is to crush 
the rock and concentrate the copper 
by mechanical processes followed by 
melting in large reverberatory fur- 
naces to separate the copper from the 
remaining gangue. If the ore is a 
sulfide it is crushed and concentrated 
by washing and flotation as described 
under zinc. The latter process is so effective that from an ore 
running as low as 2 per cent, copper 90 to 96 per cent, of the 
metal can be obtained in the concentrate. 

The first step involved in the actual smelting of the ore is the 
formation of a concentrated mixture of copper and iron sulfides 
known as copper matte. This is done by melting a properly 
proportioned mixture of the concentrated ore and limestone 
either in a reverberatory furnace, if the ore is in fine particles, or 
in a copper blast furnace if it is coarse. In either type of furnace, 


FIG. 59. 


part of the iron sulfide is oxidized to sulfur dioxide and iron 
oxide and the latter together with lime from the limestone and 
silica from the ore forms easily fusible silicates or slags, the re- 
maining iron sulfide and the copper sulfide melt together forming 
the matte and when the charge is withdrawn from the furnace, 
the matte goes to the bottom of the receiving vessel and is easily 
separated from the slag. In the reverberatory furnace, the ore 
and limestone are spread on the horizontal bed of the furnace and 
exposed to the action of the flame produced by blowing powdered 
coal into the furnace with a blast of air; while in the copper 
blast furnace the ore and flux together with coke for fuel is charged 

FIG. 60. 

into the top of the long narrow vertical furnace (Fig. 59) into 
which air is blown at the bottom. The temperature reached is 
high and the molten slag and matte are drawn off at the bottom. 
The matte from either type of furnace is next treated in a copper 
converter (Fig. 60), which is a large, approximately barrel- 
shaped vessel resting on a rotating mechanism. The con- 
verter has a thick lining of magnesia, MgO. Air at a high 
pressure is blown into the converter from several blast pipes 
on the side. 

The large supply of air, oxidizes the ferrous sulfide in the 
matte to iron oxide and sulfur dioxide, thus liberating heat which 
serves to keep the mass fluid. 



2FeS -f- 30 2 = 2FeO + 2S0 2 

The sulfur dioxide escapes as gas, while the ferrous oxide 
combines with silica which is added along with the matte to 
form ferrous silicate slag, 

FeO + SiO 


Some cuprous sulfide is oxidized, but the cuprous oxide formed 
reacts at once with iron sulfide reforming cuprous sulfide, 

2Cu 2 S + 3O 2 = 2Cu 2 + 2SO 2 
Cu 2 + FeS = Cu 2 S + FeO 

After the iron has been eliminated the slag is poured off and the 
second period of blowing begins. The cuprous sulfide is now 
oxidized to some extent, forming cuprous oxide. This reacts 
with the remaining cuprous sulfide forming sulfur dioxide and 
metallic copper 

Cu 2 S + 2Cu 2 = 6Cu + SO 2 

At the completion of this operation metallic copper remains 
and this is poured into molds. It is known as blister copper and 

















f 'W 

1 CPP 







.- : 

^ Anc 










r i 



/, -A//- 








3 fatt 

''Solution of Copper Sulfate andSu/furic Add 
FIG. 61. 

carries practically all of the nickel, gold, silver, and platinum 
that were in the original ore. It also contains some other im- 
purities and is not fit to be used for most commercial purposes. 
Further purification is best obtained by means of electrolysis. 

Electrolytic Refining. For this process, the crude blister 
copper is cast into large plates which are used as anodes in elec- 
trolytic cells (Fig. 61) where the cathodes are thin plates of pure 
copper, and the electrolyte is an acidified solution of copper 


sulfate. On passing a suitable current through the cells, 
copper is dissolved from the impure, anodes and deposited 
in nearly pure condition on the cathodes. The impurities in 
the anode are either dissolved and held in solution in the elec- 
trolyte or remain undissolved and fall to the bottom of the cell. 
Gold, silver and platinum are thus precipitated and are recovered 
by subsequently melting up the slimes or mud from the bottom 
of the cells. The electrolytic copper so obtained is nearly pure. 
It is melted again, stirred with wooden poles to remove oxide and 
then cast into bars for the market. 

Physical Properties of Copper. Copper is red by reflected 
light and is one of the very few colored metals. When a very 
thin layer of copper is viewed by transmitted light it is bluish- 
green. The density of the metal is 8.94 and it melts at 1,083 
and boils at 2,310. The metal is a very good conductor of heat 
and only silver exceeds it in conductivity for electricity. A large 
proportion of the copper which is put upon the market is used for 
electrical conductors. For this purpose it must- be extremely 
pure. It is highly malleable, ductile and tenacious, so it may 
be rolled or hammered into thin sheets or drawn into fine wire. 

Chemical Properties. Copper is not oxidized in perfectly dry 
air but in ordinary air it soon becomes coated with an adherent 
film of cuprous oxide and finally, especially if exposed to the 
weather, with a basic carbonate. The underlying metal is 
thoroughly protected by these films and is very durable. On 
this account, it is used for the roofs of important buildings and 
as a sheathing for ships. At somewhat elevated temperatures 
copper combines directly with oxygen, chlorine, and sulfur. 
It does not act on water at any temperature and is not dissolved 
by acids with evolution of hydrogen. This is connected with 
the fact that the cupric ion is a stronger oxidizing agent than 
the hydrogen ion. Such acids as contain elements in a condition 
which makes them stronger oxidizing agents than hydrogen as 
ion, nitric or hot concentrated sulfuric acid for example, will 
dissolve copper but with the reduction of the acid and not the 
evolution of hydrogen. On the other hand, when in contact 
with the air, acids like hydrochloric and even acetic dissolve the 
metal. The oxygen of the air is the oxidizing agent in this case. 

The atomic weight of copper is 63.57. 


Alloys. Copper forms many varied and useful alloys with 
the other elements. With zinc it forms brass, 30 to 37 per cent, 
zinc, and tombac 2 to 15 per cent, zinc, and with zinc and iron, 
an alloy which is malleable when hot. Bronze contains essen- 
tially copper and tin, although other metals are often added, 
especially zinc. Gun metal contains 10 per cent, tin and bell 
metal from 17 to 20 per cent. tin. Speculum metal is one-third 
tin and contains a little arsenic. It takes a very high polish. 
Phosphor bronze contains tin and a little phosphorus. Alu- 
minum bronze has from 3 to 10 per cent, of aluminum. That 
with the higher percentage of aluminum has about the color of 
gold, is easily cast, takes a high polish and is nearly as strong as 
cast steel. Manganese bronze has a content of about 30 per 
cent, manganese. Silicon bronze contains up to 5 per cent, 
silicon. It is sometimes used for exposed electric wires since 
its tenacity is greater than that of copper although its conduc- 
tivity is smaller. German silver is an alloy of copper, nickel and 
zinc and usually contains about 22 per cent, each of zinc and 

Cuprous Compounds. The cuprous or monovalent compounds 
are formed as the first step in the oxidation of the element or 
in the reduction of the cupric, or divalent compounds. In 
some ways many of them are more stable than the cupric com- 
pounds, and are formed from the latter by spontaneous decom- 
position; but in general they are relatively unstable and are 
transformed by oxidizing agents into the cupric compounds. 
The principal cuprous compounds are Cu 2 O, CuCl, CuBr, Cul, 
CuNC, and Cu 2 S. Some of the cuprous oxyacid salts can be 
obtained in solution, but they soon decompose in such a manner 
that half the cuprous is reduced to metallic copper while the 
other half is oxidized to cupric. The cuprous ion Cu + seems to 
be incapable of existing in any appreciable concentration for 
all the soluble stable salts are those of complex ions. 

When a given quantity of electricity is passed through a 
cuprous salt, just twice as much copper is deposited as when 
the same quantity is passed through a cupric solution. This 
of course, would follow from the relation between their valencies. 

Cuprous Oxide. Cuprous oxide occurs in nature and is 
known as cuprite or red oxide of copper. It may be prepared 


by the direct union of the elements, or by the reduction of an 
alkaline solution of a cupric salt with glucose. It has a red 
color and is not soluble in water. It dissolves in hydrochloric 
acid forming the compound, HCuCl 2 which gives the complex 
ion CuCl 2 ~~; and in ammonium hydroxide owing to the formation 
of the complex ion Cu(NH 3 )2 + . Both of these ions are colorless. 
Cuprous hydroxide is not stable. 

Cuprous Chloride. Cuprous chloride, CuCl, is one of the most 
important of the cuprous salts. It is most readily prepared by 
adding hydrochloric acid to a solution of cupric chloride, CuCl 2 , 
and boiling for some time with finely divided copper out of 
contact with air. The equation for the reaction is : 

CuCl 2 + Cu + 2HC1 = 2HCuCl 2 

When the solution is poured into water, the compound HCuCl 2 
breaks up, 

HCuCl 2 4=HCl + CuCl 

Cuprous chloride is white and is but slightly soluble in water. 
It is dissolved by concentrated hydrochloric acid and soluble 
chlorides in general through the formation of the complex ions 
CuCl 2 ~ or CuCl 3 . These solutions are colorless when pure, 
but quickly become colored when exposed to the air, owing to 
the oxidation to cupric compounds. Cuprous chloride is also 
soluble in ammonium hydroxide solutions owing to the forma- 
tion of the colorless ion Cu(NH 3 ) 2 + . The solution, however, 
quickly turns blue upon exposure to the air because of oxidation 
to the highly colored cupric ammonia ion. 

Both the ammoniacal and hydrochloric acid solutions of 
cuprous chloride have the property of absorbing carbon monoxide 
through the formation of an unstable compound. Advantage is 
taken of this in gas analysis. 

Cuprous Bromide. Cuprous bromide, CuBr, is very similar 
to the chloride. It may be prepared by simply heating the 
anhydrous cupric bromide out of contact with the air when it 
decomposes into bromine and the cuprous salt. The chloride 
shows a tendency in the same direction but the decomposition 
is incomplete. 

Cuprous Iodide. Cupric iodide is so unstable that it decom- 
poses at ordinary temperatures even when in solution; so when 


a soluble iodide is added to a solution of a cupric salt, cuprous 
iodide is at once precipitated and iodine liberated, 

2Cu + + + 4 r ?= 2CuI + I 2 

The cuprous iodide is colorless but appears brown from the 
color of the liberated iodine. The reaction is reversible and is 
far from complete unless the iodine is reduced by the addition 
of a sulfite or other suitable reducing agent. 

Cuprous Cyanide. When potassium cyanide is added to a 
cupric salt, cupric cyanide is precipitated; but this is unstable 
and soon decomposes into the cuprous cyanide, CuNC, and 

2Cu(NC) 2 = 2CuNC + C 2 N 2 

The cuprous cyanide is practically insoluble in water, but dis- 
solves readily in an excess of potassium cyanide owing to the 
formation of the colorless very stable complex anion Cu(NC) 2 ~~. 
This ion gives an exceedingly low concentration of copper as ion, 
so low in fact, that it does not show the properties of copper as 
ion and will not give a precipitate of copper sulfide (which 
see) with hydrogen sulfide. Because of the very small con- 
centration of the copper ion from this cuprocyanogen ion, 
all the slightly soluble salts of copper will dissolve in an excess 
of potassium cyanide. 

Cuprous Sulfide. Cuprous sulfide, Cu 2 S, is formed by heating 
the cupric compound in a stream of hydrogen. It is found in 
nature and is then known as copper glance. It is dark colored 
and is not soluble in water or cold dilute acids, but is dissolved 
by hot dilute nitric acid and by potassium cyanide. 

The Cupric Compounds. The cupric compounds are divalent 
and are in general more common and more important than the 
cuprous. The halogen compounds, as noted in the discussion 
of the cuprous salts, show a marked tendency to pass over into 
the cuprous salt and the free halogen. On the other hand, the 
only stable oxy salts of copper belong to the cupric series. When 
in dilute solution in water, the cupric salts all show a blue color 
which consequently is said to be the color of the cupric ion, 
Cu + +. When an excess of ammonium hydroxide is added to a 
cupric salt, the precipitate of cupric hydroxide or basic salt which 


is first formed is redissolved and the solution becomes intensely 
blue. Such a solution acts as though it contained a complex 
cupric ammonia ion, Cu(NH 3 )4 + + , having a deep blue color. 

Cupric Oxide and Hydroxide. Cupric oxide or black oxide of 
copper, CuO, occurs in nature in the ore melaconite. It may be 
prepared by heating the metal in a stream of oxygen or air; 
cuprous oxide is first formed. When heated to redness, it acts 
as a good oxidizing agent toward hydrogen and organic sub- 
stances, oxidizing them to water and carbon dioxide and being 
itself reduced to metallic copper. This property is employed 
in the quantitative synthesis of water and in the analysis of 
organic compounds. Copper oxide is also used in the refining 
of petroleum to remove sulfur compounds. 

Cupric hydroxide, Cu(OH) 2 , is a weak base which is but 
slightly soluble in water. It is precipitated as a light blue gelat- 
inous substance when a soluble hydroxide is added to a cupric 
salt. Upon standing or more quickly when the solution is 
boiled, it becomes black and is transformed into a hydrated 
form of the oxide. 

Cupric hydroxide is of course soluble in acids. It is not dis- 
solved by sodium hydroxide of moderate concentrations, but 
is soluble in ammonium hydroxide owing to the formation of 
the deep blue complex cupric ammonia ion Cu(NH 3 )4 + + . It is 
also soluble in solutions of sodium and potassium tartrate, espe- 
cially in the presence of sodium hydroxide, a complex copper 
tartrate ion, CuCJHUOe , is formed. This solution is deep blue 
like the ammoniacal solution, and is used in the determination 
of glucose since the latter reacts in a definite manner with it 
forming cuprous oxide which is presipitated. It is called Fehl- 
ing's solution. This solution of cupric hydroxide in ammonia 
has the remarkable property of dissolving cotton or other forms 
of cellulose. This is reprecipitated upon acidulating the solution 
in a form having a silky luster. This is the basis of the process 
of making artificial silk. 

Cupric Chloride. Cupric chloride, CuCl2, may be prepared 
by the direct union of its elements or by the interaction of hydro- 
chloric acid with the oxide, hydroxide, or carbonate. 

Anhydrous cupric chloride is yellow, a dilute solution of the 
salt is blue and a concentrated solution is green, perhaps because 



of the simultaneous presence of the ion and of the undissociated 
salt. From the. concentrated solution, the blue hydrate, 
CuCl 2 -2H20, crystallizes out. 

The deep blue solution formed by the addition of an excess 
of ammonia to a solution of cupric chloride deposits upon con- 
centration crystals having the formula Cu(NH 3 ) 4 Cl 2 -H 2 O. The 
anhydrous chloride will combine directly with ammonia to 
form the compounds, CuCl 2 -6NH 3 and CuCl 2 -2NH 3 , these are 
analogous to salts with water of crystallization. 

Cupric Sulfate. Copper sulfate or blue vitriol, CuSO 4 -5H 2 O, 
is the most important copper salt. It may be made on a small 
scale by heating copper with concentrated sulfuric acid. The 
sulfuric acid here acts as the oxidizing agent, and is itself reduced 
to sulfur dioxide, 

Cu + 2H 2 S0 4 = CuSO 4 + S0 2 + 2H 2 O 

On a large scale, it is made by carefully oxidizing naturally 
occurring sulfides by roasting them in the air and dissolving out 
the sulfate with water, or by roasting the ores or matte and dis- 
solving in dilute sulfuric acid. The salt is soluble in water, and 
crystallizes from solution as the blue pentahydrate, CuSO 4 -5H 2 0. 
The five molecules of water of crystallization in the pentahy- 
drate may be successively replaced by the same number of mole- 
cules of ammonia. A similar thing is true for other salts and 
serves to strengthen the resemblance between ammonia and 
water. With potassium sulfate, copper sulfate forms the double 
salt, CuS0 4 -K 2 SO 4 -6H 2 O, isomorphous with the corresponding 
magnesium compound. 

Copper sulfate is used in the preparation of other copper com- 
pounds, in calico printing, in the purification of copper by elec- 
trolysis, in electroplating and electrotyping, as a germicide and 
fungicide (mixed with slaked lime it forms Bordeaux mixture). 
Since it is very poisonous to the lower orders of plants, it is used 
extensively to destroy such organisms as often give to public 
water supplies a diasgreeable taste and odor. It is used in mak- 
ing the electric cell known as the Daniel's battery and also in the 
gravity battery. 

Cupric Nitrate. Copper nitrate, Cu(N0 3 ) 2 -6H 2 O is a blue 
salt which may be made by dissolving the metal, the oxide or 


carbonate in nitric acid. It is very easily soluble in water. 
When heated it gives the oxide. 

Cupric Carbonate. Cupric hydroxide is such a weak base that 
the normal carbonate does not seem to be able to exist since 
it has neither been found in nature nor prepared in the labora- 
tory. Two basic carbonates, malachite, CuC0 3 -Cu(OH) 2 and 
azurite, 2CuC0 3 -Cu(OH) 2 occur fairly abundantly in nature and 
are valuable copper ores. The green coating which gradually 
forms on copper exposed to the weather has the composition of 

Cupric Acetate. A basic copper acetate, known as verdigris 
having a fine bluish-green color and used as a pigment, is formed 
by the action of the air and of the crudest kind of vinegar (acetic 
acid) upon plates of copper. The oxygen of the air is the oxidiz- 
ing agent in this case. The normal salt, Cu(C 2 H 3 O 2 ) 2 -H 2 O may 
be obtained by crystallization from dilute acetic acid. It forms 
dark green crystals and is used as a pigment. 

Cupric Sulfide. When hydrogen sulfide is passed through a 
neutral or acid solution of a cupric salt, a brownish-black pre- 
cipitate is thrown down; this consists principally of the cupric 
sulfide, CuS, but always contains some cuprous sulfide, Cu 2 S. 
By drying this precipitate and heating it in a current of hydrogen, 
it is completely transformed into the cuprous compound. The 
hydrogen assists by sweeping away the sulfur. 

Cupric sulfide is practically insoluble in cold dilute acids, but 
dissolves in hot dilute nitric acid. In the latter case, hydrogen 
sulfide is not evolved; but the sulfur as ion is oxidized to free 
sulfur and the nitric acid reduced to nitric oxide. 

Copper Ferrocyanide, Osmotic Pressure. Copper ferro- 
cyanide, Cu2Fe(NC) 6 has a very intense brownish-red color 
and is exceedingly slightly soluble. Because of these properties, 
it is of some importance to analytical chemistry. It is, however, 
of far greater importance to biology and to theoretical chemistry, 
because of the part which it plays in the study of the phenomena 
of osmotic pressure. 

Copper ferrocyanide is a colloidal substance, and may be 
obtained as a membrane filling the pores of a porous cell (Fig. 62) 
by putting a solution of copper sulfate on the inside of the cell 
and a solution of potassium ferrocyanide on the outside and 



allowing the whole to stand for some time. The ferrocyanide is 
precipitated in the pores of the cell where the two solutions meet. 
After this cell has been thoroughly washed to free it from soluble 
salts, it behaves in a very peculiar manner. Pure water will 
pass through it in either direction very much as through an 
ordinary porous cell only more slowly; but if a solution of sugar 
or of many other substances is placed in the cell, and the latter 
then set into a dish of pure water, the water will be drawn through 
into the solution. If the cell be closed with a 
cork, the water will continue to enter until it 
has produced a certain pressure which can be 
measured upon an attached manometer. When 
this pressure is reached, the water ceases to enter, 
and if by any means the pressure within the cell 
is increased pure water and not solution is forced 
out of the cell until the pressure is reduced to 
its former value. This pressure at which there 
is equilibrium between pure water and a solution 
separated by a membrane such as copper ferro- 
cyanide is called the osmotic pressure of the 

When the conditions which affect the osmotic 
pressure are investigated, it is found that for 
fairly dilute solutions, the osmotic pressure is 
directly proportional to the concentration of the 
solution and to the absolute temperature at which 
it is measured. In these ways it corresponds 
exactly to the pressure of a gaseous substance. 
In fact, the osmotic pressure of a solution is the 
same as the pressure which the solute would have 
if it were in the state of a gas and occupied the 
volume of the solvent in which it is dissolved at the temperature at 
which the osmotic pressure is measured. The gas law equations, 
PV = K and V = K'T apply to osmotic pressure also, P is the 
osmotic pressure, V the volume of the solvent (not the volume 
of the solution), and T the absolute temperature. The osmotic 
pressure of a solution containing one gram mole in a liter of 
solvent is approximately 22.4 atmospheres at 0C. 

Other membranes than copper ferrocyanide may be used 

FIG. 62. 


without altering the value of the osmotic pressure of a solution 
provided always that they will permit the solvent to pass through, 
but not the solute. Such membranes are called semipermeable 
membranes. They seem to act as selective solvents and to 
dissolve the solvent but not the solute of the main solution. 

The osmotic pressure of salt solutions is abnormally high 
having about twice the calculated value, in the case of sodium 
chloride, for example. The significance in connection with the 
dissociation hypothesis is obvious, but really it adds nothing to 
the evidence already given in favor of this theory, because the 
science of thermodynamics has shown that the osmotic pressure 
of a solution is directly proportional to the relative lowering of the 
vapor pressure of the solution and hence to the lowering of the 
freezing-point and to the rise in the boiling-point. If these are 
abnormal, as they are, for salts, the osmotic pressure must be 
abnormal also. 

The osmotic pressures of concentrated solutions are by no 
means small and pressures up to 36 atmospheres have been 
measured. Such a pressure is equal to that exerted by a column 
of water 1,200 ft. in height. And this enormous pressure it 
must be remembered, is produced simply by the water's passing 
through a semipermeable membrane into a solution without the 
action of any external force. 

Many of the processes which take place in the living organisms 
are at least partially osmotic in their nature, the movement of sap 
in plants and of the bodily fluids in animals, for example. 

Analytical Properties of Copper. In analysis, use is most 
frequently made of the following properties: namely, the 
sulfide is but slightly soluble in water and cold dilute acids, 
but soluble in hot dilute nitric acid (separation from mercury) ; 
the sulfate is soluble in water (separation from lead); the hy- 
droxide is soluble in an excess of ammonia (separation from 
bismuth) ; and the sulfide is soluble in potassium cyanide (sepa- 
ration from cadmium). The intense blue color of the cupric 
ammonia ion which is formed by the addition of an excess of 
ammonium hydroxide to the cupric salt is fairly characteristic, 
since nickel is the only other metal which has similar properties. 
A very characteristic test for copper is based on the fact that 
metallic iron will precipitate bright metallic copper upon its 





surface when it is introduced into a solution of a copper salt. 
The formation of a precipitate of copper ferrocyanide is one 
the most sensitive tests for copper. 

In quantitative analysis, the copper is usually weighed as the 
metal, into which condition it is brought by electrolysis of its 
sulfate or nitrate with platinum electrodes. 

Voltaic Cells. Yoltaic cells are devices for the conversion of 
free 1 chemical energy into electrical. They all have the following 
characteristics in common. They are made up of spontaneously 
occurring oxidation and reduction reactions 
which are capable of taking place at two 
different points. The reactions in the 
cell must be spontaneous because these 
cells can transfer work from one system 
to another, and therefore must lose free 
energy; and only such changes as take 
place of themselves are accompanied by a 
decrease in free energy. These changes 
must be oxidation and reduction, because 
such processes take place with the transfer 
of positive electricity from the oxidizing 
agent to the reducing agent, and the transfer of electricity is an 
essential part of the voltaic phenomena. They must be able to 
take place at two different points, so that positive charges may 
be given up by the oxidizing agent to an electrode, and then flow 
through a wire or other conductor to another electrode where 
they are given to the reducing agent and so oxidize it. During 
this passage of electricity from one electrode to the other, it may 
be made to do the various kinds of work which a voltaic cell is 
called upon to perform. 

The Daniell cell furnishes a good illustration of all that has 
been said above. This cell (Fig. 63) may be made by filling a 
porous cup with a solution of zinc sulfate and putting in a zinc 
electrode. The cup is then placed in a solution of copper sulfate 
in which there is a copper electrode. When the two electrodes 

1 The free energy of a system is that part of its transferable energy which 
can be conveyed to another system in the form of actual work. It is usually 
different from the total energy which the system is able to hand over to the 
other one. The difference in the total energy transferred and the free is 
passed on in the form of heat or some other manifestation than work. 

FIG. 63. 


are connected by a wire, positive electricity flows through the 
wire from the copper to the zinc electrode. At the same time, 
metallic copper is deposited upon the copper electrode and 
metallic zinc passes from the zinc electrode into the solution 
and becomes zinc ions. The changes may be represented by the 
following equation: 

Zn + Cu++ + SO 4 = Cu + Zn++ + SO 4 

From this it may be seen that the cupric ion has oxidized the 
zinc to the zinc ion, and has been itself reduced to metallic copper. 
This change will quickly take place if a piece of zinc is placed in 
solution of copper sulfate; soon there will be a solution of zinc 
sulfate and a precipitate of metallic copper. It is then a spon- 
taneous change. But this arrangement does not yield electrical 
energy, and the energy of the reaction simply goes to heat the 
solution. This is because the oxidation and the reduction take 
place at the same point. In the cell which is described above, 
the oxidizing agent is kept from direct contact with the reducing 
agent by the porous cup, and the only way in which the copper 
ion can oxidize the zinc is by becoming metallic copper, thereby 
giving its positive charges to the copper electrode from which 
they may flow through the wire to the zinc electrode, and oxidize 
it to the zinc ion. While this electricity is flowing through the 
wire, it may be made to do work. 

The electrical energy which a voltaic cell is able to produce is 
equal to the decrease in the free energy that accompanies the 
chemical change which takes place. Now electrical energy is 
the product of two factors; the coulombs and the volts, the 
coulombs being the quantity factor and the volts the intensity. 
There is, then, the following relation between the free energy and 
the electrical energy: 

Free energy = volts X coulombs 

The unit of electrical potential, the volt, has such a magnitude 
that a volt X a coulomb is equal to a joule or 10,000,000 ergs. 

Now when a gram equivalent of any oxidizing agent is reduced, 
96,500 coulombs (see Faraday's Law, p. 112) of positive elec- 
tricity are transferred to the reducing agent. 

From the relation given above that the free energy is equal 



to the product of the volts times the coulombs, it follows that 
the potential or electromotive force of a cell is equal to the free 
energy produced during 1 grm. equivalent of change divided 
by 96,500. Expressed in the form of an equation, 

E.m.f. = F/96,500 

in which e.m.f. stands for electromotive force, and F for the 
free energy of 1 grm. equivalent of chemical change. From 
this, it follows that the potential of a cell is directly proportional 
to the decrease in the free energy accompanying the transfor- 
mation of a gram equivalent of the oxidizing and reducing agents 
in the cell. Since this free energy is different for different com- 
binations, the potentials of various cells are quite dissimilar. 
Oxidizing agents can be arranged in the order of their activity, 
i.e., the decrease in the free energy involved in the transforma- 
tion of a gram equivalent, and a number can be assigned to 
each which represents the potential of an electrode in the normal 
ionic solution 1 of the oxidizing agent. This number may be 
represented by E and then, E = F'/96,500; where F' is the free 
energy produced when 1 grm. equivalent of the oxidizing agent 
is reduced. The following table gives a short list of oxidizing 
agents arranged in the order of their oxidizing power, and also 
gives the value of E. This table is a part of what is called the 
potential series. 


E in volts 


E in volts 

NO 3 -. . 

+ 1 75 

H + . 

+ 0.277 

C1 2 

+ 1 694 

Pb + + 

+ 129 

Br 2 . ... 

+ 1.396 
+ 1 270 

Ni ++ 
Co + + 

+ 0.049 
+ . 045 

Ag + 

+ 1 048 

Tl + 

- 045 

Cu ++ 

+ 1 . 027 
+ . 606 

Fe ++ 
Cd + + 

- 0.143 

Zn + + 

. 493 

A + sign before the value of E indicates that the electrode is 

1 A normal ionic solution is one which contains in a liter such a quantity 
of the salt that there will be 1 grm. equivalent of the ion present in the 
actually dissociated state. Because of the incomplete dissociation, such a 
solution will contain per liter more than one equivalent of the salt. 


positive against the solution of that oxidizing agent, and a 
sign means that the electrode is negative toward the solution, 
and the significance of a negative sign is that free energy is taken 
up instead of being given out when the ion is changed to the 
metal. This is the case with all the metals which are hard to 
reduce from their compounds. 

Any oxidizing agent in this list is able to form anything 
lying below it and is formed by all that stand above. For 
example, the cupric ion will oxidize anything from hydrogen to 
zinc and change it from the element to the ion, while copper is 
itself oxidized by mercury and all that stand above to the cupric 
ion. A voltaic cell then may be made up of a combination of any 
two of these agents, and it's electromotive force will be the 
algebraical difference between the value of E for the two elec- 
trodes. For example, a cell might be made up of a silver 
electrode in a silver solution and a copper electrode in a copper 
solution, and it would have an e.m.f. of 1.048 - 0.606 = 0.442 
volts. The silver electrode would be positive toward the copper 
and the latter would dissolve, while metallic silver would be 
deposited. The electromotive force of the Daniell cell may be 
calculated from this table by subtracting the value of zinc, 
-0.493 from that of copper, +0.606; 0.606 - ( -0.493) = 
0.606 + 0.493 = 1.099 volts. Herein contrast to the above, the 
copper is positive toward the zinc and is deposited while the 
zinc dissolves. 

The activity of an oxidizing agent depends upon the concen- 
tration. The figures given above apply to normal ionic solu- 
tions containing 1 grm. equivalent of "actual ion " per liter; in 
more dilute solutions, the electrodes are less positive if the oxidiz- 
ing agent is a cation. Anything which tends to reduce the con- 
centration of the copper ion for instance, will tend to bring the 
potential of the copper electrode nearer the zinc and reduce the 
potential of the copper zinc cell. The concentration of copper in 
a solution containing an excess of ammonia is very small, and 
the potential of a cell made up of a copper electrode in such a 
solution and a zinc electrode in zinc sulfate solution is much 
smaller than that of the Daniell cell. If potassium cyanide is 
used instead of the ammonia, the concentration of the copper 
is still further reduced, so that the potential of the copper elec- 


trode is actually made smaller than that of the zinc; and in a 
cell made up of the two, the processes which take place in the 
Daniell cell are reversed, copper going into solution and the zinc 
being deposited. 


General. Because of its occurrence free in nature, silver has 
been known and used from prehistoric times. Its relationship 
to the other elements is rather complex as will be seen by a study 
of its compounds. Physically it is much like copper aside from 
the difference in color; chemically it is like mercury in the 
mercurous, and copper in the cuprous states in that it is mono- 
valent and also in the slight solubilities of its halogen compounds. 
It differs radically though from both mercury and copper in that 
its hydroxide is a strong base, and that' its soluble salts of strong 
acids are neutral. In this it resembles the alkali metals, and some 
of its salts are isomorphous with the corresponding salts of potas- 
sium. The silver ion Ag + is a good oxidizing agent, and the 
metal is obtained very easily from many of its compounds. The 
oxide for example, when heated, acts like mercuric oxide in that 
it is changed into the metal and oxygen. Scheele took advantage 
of this in one of his methods for the preparation of oxygen. 

Occurrence. Like copper, silver is found in nature both free 
and in combination. The native silver is a fairly important 
ore, but is practically never found pure, being alloyed with gold, 
copper or mercury. The most important ore is the sulfide Ag 2 S 
called argentite or silver-glance. This is usually found in iso- 
morphous mixture with lead sulfide. In addition there are com- 
pounds of silver sulfide and antimony or arsenic sulfides known 
as pyrargyrite Ag 3 SbS 3 , and proustite, AgsAsSs. 

Metallurgy. The metallurgy of silver is complicated by the 
fact that it is generally associated with other metals such as 
copper, lead, zinc and gold which it is necessary and desirable 
to extract at the same time. To meet the needs of individual 
cases, a considerable variety of methods for extracting the 
silver have been devised. These may be grouped into smelting 
or dry processes, amalgamating and leaching or wet processes. 
The smelting methods are usually applied to ore which contains 
relatively more copper or lead than silver. If they are mainly 


copper ores, they are worked as described, under that metal, 
and the blister copper obtained will parry practically all the 
silver in the ore. This crude copper is then purified by electroly- 
sis, and the silver is obtained in the anode mud. The lead 
ores are smelted as will be described under that metal, and the 
lead will contain the silver. This is called work lead or base 
bullion. The silver is removed from this by what is known 
as the Parke's process which is as follows: Work lead is melted 
in large kettles and kept molten for some time to allow other 
impurities to be oxidized and skimmed off as "dross." Then 
a quantity of zinc equal to from 0.5 to 2.0 per cent, of the weight 
of the lead is stirred into the molten mass. Zinc and lead 
are only slightly soluble in each other, while zinc is a much 
better solvent for silver than lead is. The zinc, therefore, 
acts as an extractive solvent, and dissolves the greater part 
of the silver (see the Law of Distribution) from the lead. The 
silver zinc alloy being much the lighter, rises to the surface, and 
on being allowed to cool, forms a solid crust while the lead is 
still molten. This is skimmed off and freed from most of the 
lead which adheres to it, and the zinc is then distilled from the 
alloy in large graphite retorts. The silver is heated in a cupella- 
tion furnace where any lead which it retains is oxidized to 
litharge, PbO, and absorbed in the bone ash which forms the 
bottom of the furnace. The silver remains unoxidized, accom- 
panied by any gold that was in the original bullion. These 
are separated by dissolving the silver in hot concentrated sul- 
furic acid; the gold is left unchanged. The silver is precipi- 
tated from the silver sulf ate solution by means of metallic copper. 

There are also electrolytic processes for getting silver and 
gold from work lead. The silver and gold are found in the 
residues in the bottom of the cell. 

Amalgamation Process. This depends upon the solution 
of free silver from an ore in metallic mercury, forming silver amal- 
gam. After separating this from the ore gangue, it is freed from 
excess mercury by straining through canvas or buckskin. The 
solid amalgam is then distilled in an iron retort, the mercury pass- 
ing over and the silver remaining in the retort. This process is 
applicable to ore in which the silver is either free or can be easily 
reduced to the free state. 


Leaching Processes. In these processes, silver contained in 
the ores or metallurgical products is first converted into com- 
pounds soluble in water or in solutions of certain reagents such 
as strong brine, potassium or sodium cyanide, or sodium thio- 
sulfate in water; these compounds are then dissolved or leached 
out, and the silver precipitated from the solution by suitable 

The annual production of silver in the United States is about 
66,000,000 oz. worth about $40,000,000. 

Physical Properties. Pure silver has a very high luster and a 
beautiful white color. It melts at 955 in the presence of air or 
at 960 in its absence. The difference is due to the fact that 
oxygen is soluble in molten silver, 22 volumes of oxygen to 1 of 
silver, and this of course lowers the melting-point of the metal. 
When a molten mass of silver cools, the outside hardens before 
the liquid interior has lost its oxygen, and the latter forcing its 
way out as the cooling continues breaks the crust of solid silver 
and ejects portions of the metal from the interior producing very 
strange shapes. The phenomenon is known as " sprouting." 

Silver has a density of 10.53 and boils at 1,955. It is very mal- 
leable and ductile so that it may be drawn into wire so fine that 
a mile of it will weigh less than a gram. 

When an electric arc is formed between silver wires under 
water, a colloidal solution of silver is produced. Similar colloidal 
solutions may be made by the action of certain reducing agents 
upon silver solutions. These solutions vary in their color, 
brown, white, blue, red, and yellow being well known. 

Because of its high luster, beautiful color, and chemical 
resistivity, silver is much used for ornaments and for coinage. 
For such purposes, the pure metal is too soft so it is alloyed 
with copper. Silver coins usually contain 10 per cent, of copper. 
The coinage of England, however, contains 7.5 per cent, of 
copper and this is also the composition of "sterling silver." 

Silver is the best conductor of heat and of electricity that is 

Chemical Properties of Silver. Under ordinary conditions, 
silver does not combine with oxygen at any temperature, but it is 
oxidized at ordinary temperatures by moist ozone forming silver 
peroxide, Ag202- It will combine slowly with sulfur and the 


halogens. Because of its resistance toward oxygen at high 
temperatures, silver was formerly called one of the " noble 

It is even harder to oxidize silver to the ionic state than it is 
copper, and consequently it does not dissolve in the ordinary 
acids with the evolution of hydrogen. It is, however, oxidized 
and dissolved by nitric acid and by hot concentrated sulfuric 
acid with the evolution of oxides of nitrogen or of sulfur, the 
reduction products of these acids. Silver sulfide is an excep- 
tionally stable substance, and its formation takes place with a 
great decrease in free energy; because of this, hydrogen sulfide 
is able to react with silver with the evolution of hydrogen. 

Silver is not attacked by the alkalies either in solution or in 
the fused state, so silver dishes are often used in working with 
these substances. Platinum and other metals whose oxides have 
a tendency to form acids are attacked by strong bases especially 
in the fused state. 

The atomic weight of silver is unusually important since the 
atomic weights of a number of other elements are determined 
with its aid, it is 107.88. 

The Oxides of Silver. Silver oxide is deposited as a brown 
precipitate which darkens on drying, when a solution of sodium 
or potassium hydroxide is added to a solution of a silver salt or 
is boiled with silver chloride. The oxide is slightly soluble in 
water and the solution has an alkaline reaction showing that it 
contains silver hydroxide, AgOH, but this evidently breaks 
down into silver oxide and water. In spite of this, silver hydrox- 
ide is a strong base and its solutions take up carbon dioxide and 
precipitate the carbonate. It dissolves in ammonium hydroxide 
because of the formation of the silver ammonia ion; the result- 
ing solution contains silver ammonium hydroxide, Ag(NH 3 ) 2 OH, 
which is a strong base. The oxide is dissolved by acids giving 
solutions of the corresponding salts. 

Silver oxide is reduced by hydrogen at temperatures as 
low as 100, and when heated alone it decomposes at 250 
to 270. 

Silver peroxide, Ag 2 2 , is formed by the action of moist 
ozone upon the metal; it is unstable. A suboxide Ag 4 0, is 


The Halogen Compounds of Silver. Silver fluoride may be 
made by the action of hydrofluoric acid upon the oxide or car- 
bonate. It is highly soluble in water and is deliquescent. 

The chloride, bromide, and iodide of silver are very slightly 
soluble in water or in dilute acids, and are thrown down as 
curdy precipitates upon the addition of the corresponding 
halogen salts to a solution of a silver salt. The chloride is white 
and soluble to the extent of 0.002 grm. per liter of water. The 
bromide has a pale yellow color and is less soluble than the 
chloride. The iodide is distinctly yellow, and is the least soluble 
of the three. When exposed to the light, silver chloride and bro- 
mide darken and some free halogen is formed together with a sub- 
chloride or bromide, Ag 2 Cl or Ag 2 Br. Silver iodide is unchanged 
in the light unless an excess of AgNO 3 is present. These facts 
are important in connection with photography (p. 384). 

The silver halides at ordinary or lower temperatures will 
combine directly with ammonia for the formation of the following 
compounds, 2AgCl-3NH 3 , AgCl-3NH 3 , AgBr-NH 3 , 2AgBr-3NH 3 , 
AgBr-3NH 3 , 2AgI-NH 8| Agl-NH,. 

Silver chloride is fairly soluble in concentrated solutions of 
hydrochloric acid and of the easily soluble chlorides, apparently 
because of the formation of the ion AgCl 3 ~~. It is also easily 
soluble in ammonium hydroxide owing to the formation of the 
stable complex ion, Ag(NH 3 ) 2 + called the silver ammonia ion. 
Silver bromide is somewhat soluble in ammonium hydroxide, 
but the iodide is scarcely affected. The explanation is that silver 
bromide because of its slight solubility gives a concentration of 
silver as ion in the solution which is but little higher than that 
of the silver ion which together with the ammonia is in equilib- 
rium with the complex silver ammonia ion, so its concentration 
cannot be much reduced through the formation of this ion. 
Silver iodide is still less soluble, and the concentration of the 
silver ion in its solution is smaller than that from the silver 
ammonia ion. 

Sodium thiosulfate interacts with silver salts for the forma- 
tion of the complex crystalline salt, 2NaAgS 2 3 -Na 2 S 2 3 . When 
in solution, this seems to give the complex ion, AgS 2 3 ~. The 
concentration of the silver as ion from this seems to be about 
the same as that from silver iodide for sodium thiosulfate does 


not have much effect upon the latter although it dissolves silver 
bromide with ease. 

Silver Cyanide. Soluble cyanides give with the soluble silver 
salts a precipitate of silver cyanide which is somewhat less 
soluble than the chloride. The cyanide is soluble in ammonium 
hydroxide and very easily in an excess of potassium cyanide 
owing to the formation of the very stable complex silver cyano- 
gen ion, Ag(NC) 2 ~, similar to the complex cuprocyanogen ion, 
Cu (NC) 2 ~. The concentration of the silver ion from this ion is 
so very small that silver iodide, and all of the other difficultly 
soluble salts of silver, are dissolved by potassium cyanide in 
excess. The sulfide requires the largest excess. 

Solutions of this potassium silver cyanide are used in silver 
electroplating. The object to be plated is made the cathode 
and a plate of pure silver the anode. 

Silver Nitrate. Silver nitrate, AgNOs, is the most important 
salt of silver. It is made by the action of nitric acid upon the 
metal ; 

3Ag + 4HNO 3 = 3AgNO 3 + NO + 2H 2 O 

Silver nitrate is a white crystalline easily soluble salt, iso- 
morphous with potassium nitrate. It melts at a temperature 
a little over 200 and when cast in sticks is sometimes used in 
medicine as a caustic under the name of lunar caustic. Its use as 
a caustic depends upon the property of rendering insoluble certain 
nitrogenous substances, called the albuminoids. 

In the pure state, it is not altered by light; but in the presence 
of reducing agents, it is changed to finely divided metallic silver 
having a black color. This is the cause of the stains which 
it produces on the skin. The nitrate is much used in the labora- 
tory for the detection and estimation of the halogens and is 
also the starting-point for the preparation of most of the other 
silver salts. 

Silver Sulfate. Silver sulfate may be made by the action of 
hot concentrated sulfuric acid upon the metal; 

2Ag + 2H 2 S0 4 = Ag 2 S0 4 + S0 2 + 2H 2 O 

Advantage is taken of this reaction in the " parting" of the alloys 
of gold and silver which are obtained in the smelting and refining 


of these metals. The silver is dissolved while the gold is not. 
It is moderately soluble in water and will combine with aluminum 
sulfate forming the alum AgAl(SO 4 )2.12H 2 0, which is isomor- 
phous with the corresponding potassium salt and with the whole 
series of alums. 

Silver Sulfide. Silver sulfide, Ag 2 S, is so very slightly soluble 
that it is formed by the action of hydrogen sulfide upon all 
silver compounds except the complex potassium silver cyanide 
in the presence of a very large excess of potassium cyanide. 
It is even formed, as noted above, by the action of hydrogen 
sulfide or of other sulfides upon metallic silver. 

Silver Carbonate. Silver carbonate, Ag 2 CO 3 , is a pale yellow 
salt slightly soluble in water to which it imparts a faint alkaline 
reaction. It resembles lithium carbonate in that it is more 
soluble in the presence of carbon dioxide. The fact that the 
normal salt exists indicates that silver hydroxide is a strong base. 

Other Salts. Silver thiocyanate, AgNCS (white), silver chro- 
mate, Ag 2 Cr04 (red), silver arsenite, Ag 3 AsO 3 (yellow) and 
silver arsenate, Ag 3 As0 4 (reddish-brown) are a few of the salts 
which are used in analysis. 

Photography. The most important of the modern photo- 
graphic processes depend upon the sensitiveness of the silver 
halides toward light. The "dry plate " upon which the picture 
is originally taken consists of a plate of glass or a film of celluloid 
covered with a dried emulsion of silver bromide in gelatine. 

When such a plate is exposed for a fraction of a second to light, 
no visible action takes place; but if it is then placed in a suitable 
reducing agent called the " developer," the portions of the silver 
bromide which were exposed to the light are reduced to metallic 
silver at a rate which is proportional to the intensity of the light 
and the time of exposure. The reduced silver is deposited as a 
black substance in the very position of the bromide from which 
it was obtained. The unexposed bromide will also be reduced by 
the developer but very much more slowly than that which has 
been acted upon by light. If the exposure was made in a camera, 
there will be developed upon the plate an image which in its light 
and shade effects is the reverse of the object, that is, it will be 
dark where the object is light. This is called the negative. 
After the plate is "developed," it has then to be " fixed;" that 


is to say, the unchanged silver bromide must be removed so 
that the plate may not be further acted upon by the light. This 
is done by taking advantage of the solubility of the bromide in 
sodium thiosulfate or "hypo" as it is often called. 

Just why the bromide which has been exposed to light is more 
easily reduced than the unexposed is not known with certainty, 
but it may be connected with the fact that light will transform 
it into a subbromide as noted above. 

From the negative, "positives" or "prints" are made by allow- 
ing light to pass through the negative and strike sensitized paper 
which is placed in direct contact with the gelatine film of the 
negative. The dark portions of the negative which corre- 
spond to the light portions of the object protect the paper from 
the action of light, and when the print is finished it will be light 
just where the object was. Some printing-out papers are sen- 
sitized with silver bromide and these require only very brief 
exposure, but must be developed like plates. Others contain 
silver chloride and the printing is continued until the image 
is plainly visible upon the paper. Such prints are then "toned" 
by treating them with a solution of sodium chloraurate, NaAuCl 4 , 
sodium gold chloride as it is called. The silver which has been 
reduced by the light, in turn reduces the ^pld to metallic gold 
which in the finely divided state has a pleasing reddish tint. By 
replacing the gold solution by potassium chlorplatinite, E^PtCU, 
platinum will be precipitated instead of gold, giving a very dark 
tone. By whatever process made, the prints after development 
or toning must be fixed by dissolving out the unchanged silver 
salts with sodium thiosulfate. 

Mirrors. Mirrors are now usually made by coating glass with 
silver. This is done by the action of reducing agents such as the 
tartrates, glycerine, formaldehyde, or a reducing sugar upon 
ammoniacal solutions of silver nitrate. These mirrors are far 
superior to the old-fashioned ones which were backed with an 
amalgam of tin. 

Analytical Properties of Silver. The silver ion is colorless, and 
forms a difficultly soluble compound with the chlorine ion, a 
property which is also possessed by the mercurous, cuprous, and 
lead ions. Analytically, the cuprous compounds are unimpor- 
tant. Silver chloride is not soluble in hot water, a distinction 



from lead chloride; but is soluble in ammonium hydroxide, a 
distinction from mercurous chloride. In quantitative analysis, 
silver is usually weighed as the chloride, but often as the metal, 
which is obtained either by electrolysis or by oxidizing all the 
other metals, except gold and platinum, in the air at a high 
temperature and weighing the metallic bead which is so obtained, 
then dissolving out the silver with nitric acid or sulfuric acid and 
weighing the residue of gold and platinum. The difference in 
weight is the weight of the silver. 


Gold has been known from prehistoric times. From its strik- 
ing appearance and its occurrence free in nature it must have 
been among the earliest of the metals to attract the attention 
of men. 

Gold forms a comparatively small number of compounds 
since all of its simple salts with the oxy acids are very unstable. 
In fact, gold seems to be practically incapable of existing as a 
simple ion although several complex ions are known. The 
compounds of gold belong to two series, the aurous in which 
the gold is monovalent and the auric in it is trivalent. These 
compounds are characterized by the great ease with which they 
decompose to give metallic gold. This, in general, may be 
brought about by comparatively feeble reducing agents or by a 
moderate rise in temperature. The simple aurous halogen com- 
pounds are not soluble in water and this establishes a sort of 
relationship with copper and silver. 

Occurrence. Gold generally occurs free in nature, its only 
native compounds being those with tellurium. Native gold is 
rarely found pure, being almost always alloyed with silver. It 
is carried in finely divided conditions in quartz and other rock 
materials and in pyrite, galena and other sulfides. It may be 
found in its original deposits in veins running through the main 
rock masses or in alluvial beds of sand or gravel formed by the 
breaking down of the original deposits through weathering 
action. Gold in deposits of the former kind is called reef gold 
while that of the latter kind is called placer gold. 

Among the minerals containing gold and tellurium, sylvanite, 
(AuAg)Te2 and calaverite, AuTe 2 , may be mentioned. These 


minerals are much less frequently found than native gold, but 
are of importance in some gold districts. 

Metallurgy. The processes for extraction of gold may be 
classified like those for silver as smelting, amalgamation, and 
leaching processes; but for gold there is an additional class, 
simple washing processes. 

In the latter, gold in placer deposits of sand or gravel is con- 
centrated by shaking up the material with water and allowing 
the gold to separate by its greater density from the lighter sands. 
These methods do not save all of the gold even when carried 
out under good conditions, and hence are used chiefly by pros- 
pectors, and in the opening up of new gold fields. They are 
greatly increased in efficiency if mercury is used to collect the 
concentrated gold as an amalgam and in this form they are 
rather widely used. 

Amalgamation Processes. Amalgamation processes for gold 
are very similar to those used for silver, and gold and silver are 
frequently recovered in the same operation. If the gold is dis- 
seminated in quartz or other massive rock the metal is recovered 
by crushing the ore in stamp mills and leading the crushed 
material over amalgamated plates, i.e., copper plates coated with 
mercury. The gold is held by this mercury, and from time to 
time the gold amalgam is scraped from the plate; excess mercury 
is removed by filtering the amalgam through canvas or buckskin 
and the resulting solid amalgam is distilled to free the gold from 

Smelting Processes. Gold in copper or lead ores is concen- 
trated during the smelting of these ores into the blister copper 
or the work lead. It is removed from these products, along with 
silver by the refining processes already described under silver and 

When both gold and silver are thus obtained, a separation or 
"parting" is necessary. A common method of doing this is 
to treat the gold-silver alloy with boiling concentrated sulfuric 
acid. Silver dissolves almost completely while gold is unaffected. 
The gold, after removal of silver sulfate solution, is melted, 
refined and cast into bars. Electrolytic processes are also used. 

Leaching Processes. The most important wet methods for 
gold extraction are the chlorination and the cyanide processes. 


In the former, finely crushed gold ore containing the free 
metal is treated with a solution of chlorine in water. This dis- 
solves gold, forming gold chloride, AuCls. From the solution, 
gold is precipitated as the sulfide by means of hydrogen sulfide. 

The cyanide process, patented in 1890 by Me Arthur and For- 
rest, is very widely used for the treatment of low grade ores that 
cannot be profitably worked by smelting. It is carried out by 
allowing the crushed ore to remain in contact for a considerable 
time with a very dilute solution of potassium cyanide, KNC. 
The solution dissolves both gold and silver from the ore forming 
potassium aurocyanide, KAu(NC) 2 , and potassium argenticya- 
nide, KAg(NC)2. The equations representing the reactions are 
discussed in detail under the complex cyanides of gold (p. 390). 
Oxygen is essential, for the reaction. 

To recover the precious metals the cyanide solution is passed 
through boxes containing zinc dust or turnings. Zinc dissolves 
while the gold and silver are precipitated; 

2KAu(NC) 2 + Zn = K 2 Zn(NC) 4 + 2Au 
2KAg(NC) 2 + Zn = K 2 Zn(NC) 4 + 2Ag 

Enormous quantities of gold have been extracted by the 
cyanide process from the low-grade ores of South Africa and 
this has been an important factor in cheapening gold, and hence 
raising the price of other things, and thus has contributed to the 
"high cost of living." 

The gold production for the United States in 1915 was $98,- 
891,100 which is a little more than one-fifth of the world's 

Physical Properties. Gold has a high luster and a beautiful 
yellow color. It melts at 1,062 and has a density of 19.4. 
It is exceedingly malleable and ductile and may be beaten 
out into leaves Mstbooo of an inch in thickness. Very finely 
divided gold usually has a purple color, although red, blue and 
green colloidal solutions of the metal have been obtained. Ruby 
glass owes its color to finely divided gold, as does " purple of 
Cassius" which appears to be minute particles of gold intimately 
mixed with stannic oxide. 

Gold is a very good conductor of heat and electricity being ex- 
ceeded by only silver and copper. 


Chemical Properties. Gold is even more resistant to chemical 
action than silver. It is not altered ,by oxygen at ordinary 
temperatures nor by hydrogen sulfide. It is not dissolved by 
any dilute acid nor by concentrated nitric or sulfuric acids, in- 
dicating that it requires a very strong oxidizing agent to cause 
the gold to pass into the ionic state. This is confirmed by the 
fact that it is dissolved by selenic acid which is a most active 
oxidizing agent. The metal is also dissolved by "aqua regia" 
a mixture of nitric and hydrochloric acids. 

This mixture evolves chlorine and the gold is transformed by 
it into chlorauric acid, HAuCl 4 , which gives the complex ion, 
AuCl 4 . The chlorine, of course, cannot be a stronger oxidizing 
agent than the nitric acid which set it free, but it has the advan- 
tage in this case that the product of the reaction is very stable. 
Gold is also attacked by the fused hydroxides and nitrates of 
the alkali metals forming aurates. The soluble chlorine com- 
pounds of gold are easily reduced by even such feeble reducing 
agents as the ferrous salts or oxalic acid. 

Because of its chemical resistivity, its beauty and its compara- 
tive rarity, gold has long been used for ornaments and for coinage. 
For such purposes, the pure metal is too soft and it is alloyed with 
copper or silver, usually the former. The gold coins of most 
countries contain 90 per cent, of gold or are 900 "fine"; those 
of England, however, contain 91.667 per cent, or are 916.67 
fine. The composition of gold alloys is often expressed in carats, 
pure gold being 24 carat and 18 carat gold being *%4 gold. 

The Oxides and Hydroxides. Auric hydroxide Au(OH) 3 , is 
formed when sodium hydroxide is added to an auric solution. 
It acts both as a weak base and as a weak acid. When the 
hydroxide is carefully heated, it goes over into auric oxide, 
Au 2 O 3 , which at higher temperatures decomposes into the 

Aurous hydroxide, AuOH, is formed by the action of dilute 
potassium hydroxide upon aurous chloride. When heated to 
200 it is changed into the oxide, Au 2 0, and this decomposes at 

The Halogen Compounds of Gold. Gold will combine directly 
with chlorine or bromine to form auric chloride or bromide, 
AuCl 3 , AuBr 3 . These compounds combine with hydrochloric or 


hydrobromic acid or with the corresponding halide salts, forming 
chlorauric or bromauric acids or their salts. The sodium chlor- 
aurate, NaAuCl 4 -2H 2 O, is a common commercial salt and is used 
in photography as noted above in the toning of silver prints. 
Chlorauric acid is formed by dissolving gold in aqua regia. 

When either the simple chloride or bromide or the complex 
acids are cautiously heated, aurous chloride or bromide is 
formed. At still higher temperatures, these decompose into the 
free halogen and the metal. They are not soluble in water, 
but slowly decompose in contact with it forming the auric com- 
pound and free gold. This is much like the behavior of some 
of the cuprous compounds. Auric iodide is decidedly unstable 
and passes easily into the aurous compound even at ordinary 

Sulfides. Hydrogen sulfide precipitates a mixture of aurous 
sulfide, Au 2 S, and auric sulfide, Au 2 Ss from auric solutions. 
These sulfides are soluble in the alkaline sulfides owing to the 
formation of the thioaurites, KsAuS 2 , and thioaurates, KAuS 2 . 
In this respect gold is like antimony, arsenic, and tin. 

The Complex Cyanides. One of the most important chemical 
properties of gold is its ability to form complex ions with cyano- 
gen; aurocyanogen, Au(NC) 2 ~, and auricyanogen, Au(NC) 4 ~, 
are well-known examples of such ions. They are formed by the 
action of an excess of sodium or potassium cyanide upon the 
compounds of gold. They are very stable and the solutions are 
utilized in gold electroplating, an anode of gold being used and 
the object to be plated being made the cathode. 

The aurocyanides are of especial importance, because it is 
through their formation that a large part of the gold is extracted 
from its ores. This depends upon the fact that finely divided 
gold is dissolved by solutions of potassium or sodium cyanides 
when in contact with the air. The equation for the reaction is 
as follows: 

4Au + 8KNC + 2 + 2H 2 O = 4KAu(NC) 2 + 4KOH 

The reaction is really ionic so that it succeeds as well with 
the cheaper sodium cyanide as with the more expensive potassium 



Analytical Properties of Gold. Gold may be detected by its 
reduction to the metallic state by ferrous salts, oxjaic acid, or 
stannous chloride. It is almost invariably determined by the 
fire assaying processes. This is often done as follows : a weighed 
quantity of the material containing gold is mixed with a relatively 
large quantity of lead and a little borax and heated in a shallow 
clay dish called a scorifier (Fig. 64) to a fairly high temperature in 
contact with the air. A part of the lead and practically all of 
the copper and other impurities in the gold save the silver are 

FIG. 64. 

oxidized, and these together with the "gangue" form a sort of a 
glass with the borax and the lead oxide which is formed. At the 
close of this operation the molten mass is poured into a mold and 
allowed to cool. A lead button containing the gold and silver is 
obtained. This is then heated to a high temperature in contact 
with the air in a "cupel" or little cup made of bone ash, the lead 
oxidizes and the oxide is absorbed in the cupel. There is finally 
left a button of gold and silver. The silver is then removed by 
treating the alloy with nitric acid which does not dissolve the gold. 
The fire assay is a smelting process carried out on a small scale. 



Rare earths Aluminum sub-group 

The elements of this group fall naturally into two sub-groups, 
the rare earth metals and the aluminum sub-group. The 
members of the first are scandium, yttrium, and lanthanum, 
those of the aluminum sub-group are boron, aluminum, gallium, 
indium, and thallium. The elements are all trivalent although 
some have other valencies in addition. As was the case in 
groups I and II the members which stand in the left hand column 
of the group as it appears in the periodic system, scandium, 
yttrium, and lanthanum are stronger bases than those of the 
aluminum group which are on the right. Aluminum is the third 
element in the order of abundance in the earth's crust, boron is 
fairly abundant and all the other members of the group are 
decidedly rare. 

The Rare Earth Metals. The metals of the rare earths are 
found in a large number of rare minerals each of which contains 
a great many elements. The names of a few of these minerals 
are euxenite, gadolinite, orthite, and monazite. Monazite is 
found in considerable quantities in North and South Carolina, 
but comes chiefly from Brazil. The hydroxides of the members 
of this sub-group are fairly strong bases and are not soluble in 
excess of sodium or potassium hydroxides. 


Scandium whose atomic weight is 44.1 is interesting be- 
cause it is one of the elements whose existence and properties 
were predicted by Mendelejeff. See p. 272. Its hydroxide 
Sc(OH) 3 is a colorless gelatinous precipitate soluble in acids but 
not in bases. The salts are colorless and are not especially 



hydrolyzed. The salts of the oxy-acids as well as of the halogen 
acids are known. 


Yttrium has an atomic weight of 89. The metal decom- 
poses water slowly, its hydroxide is colorless and gives colorless 
salts. It is a stronger base than scandium hydroxide. 


Lanthanum, La, atomic weight 139, is an iron-gray colored 
metal, which is very slowly acted upon by water. It dissolves 
in dilute acids. The hydroxide is white and somewhat soluble 
in water, the solution is alkaline, absorbs carbon dioxide and 
decomposes ammonium salts. It is the strongest base of the 

There are several elements, some of them of rather doubtful 
individuality, whose atomic weights are near to that of lanthanum 
and which so far as that is concerned have a claim on the same 
position in the periodic system. They have so nearly the same 
properties that it is difficult to separate them. Ostwald has 
likened this group of closely related elements with about the 
same atomic weight which appear in the periodic system where 
one element might be expected, to the planetoids which are 
found in the solar system where one large planet might be 
looked for. 

The Aluminum Sub-group. The first member of Group III 
and of this sub-group is boron, atomic weight 11. This is almost 
exclusively an acid-forming element; that is to say, it does not 
form cations, and for that reason has already been treated among 
the non-metals. 


General. Aluminum is the most abundant of the metals. 
It is trivalent, and is like zinc in that its hydroxide is soluble 
both in acids and in bases. In addition to its fairly close rela- 
tionship to the other members of this sub-group, it is very much 
like ferric iron and trivalent chromium as will be seen when these 
metals are studied. 



Occurrence. Aluminum is never found free but its com- 
pounds are exceedingly abundant, since it is an essential constitu- 
ent of nearly all rocks except the sand-stones and limestones 
and is almost invariably present in these. It is found as 
the double silicates, the feldspars and the micas; as the simple 
silicate clay; as the double fluoride with sodium in cryolite; as 
the oxide corundum; as the hydroxide and the partially dehy- 
drated hydroxide in bauxite; and in many other compounds. 
The ruby is aluminum oxide colored by a little chromium, the 
garnet is calcium aluminum silicate, and turquoise is a basic 

Preparation of the Metal. Metallic aluminum is prepared by 
the electrolysis of highly purified aluminum oxide dissolved in a 


Aluminum Oxide 
1 dissolved in 
Fused Cryo/ife 

FIG. 65. 

molten bath the basis of which is cryolite, NasAlFe. The 
electrolysis is carried out in iron cells (Fig. 65) lined with carbon 
which is made the cathode, the anodes being large carbon rods. 
The current is high, about 10,000 amp. per cell, while the drop 
of potential around the cell is about 5 volts. TJie bath is kept 
melted by the heat produced by the current in overcoming the 
resistance. Upon electrolysis after the addition of the oxide, 
the metal collects at the cathode and runs down to the bottom 
of the cell while oxygen is evolved, and combines largely with the 
carbon anodes. As the process goes on more aluminum oxide is 
added from time to time, and the metal is tapped off so that the 
operation is continuous. This process differs materially from 
most metallurgical operations in that it is impracticable to purify 
the product, and so it is necessary to start with the very purest 


raw material. The aluminum oxide is made from bauxite, and 
the simplest process for its purification is to melt the bauxite in an 
electric furnace with the proper amount of carbon. This reduces 
the oxides of iron, silicon, and titanium which are present and 
leaves the aluminum oxide untouched. The metallic alumi- 
num is very pure, often running 99.9 per cent. 

Physical Properties. Aluminum is a white metal with a bluish 
cast. It melts at 658 and has a density of 2.7 becoming 2.72 if it 
is hammered or rolled. The cast aluminum is about as hard as 
pure silver while the hammered is as hard as soft iron. Its tensile 
strength is high although much lower than that of steel. It is 
malleable and ductile and may be rolled or hammered into very 
thin sheets or drawn into fine wire. It is best worked at 100 to 
150. At temperatures near its melting-point it becomes so 
brittle that it may be easily powdered iri a mortar. It is not 
readily worked in a lathe, but many of its alloys are. 

Alloys with magnesium known as magnalium are very light, 
easily machined and fairly strong. Aluminum bronze has al- 
ready been discussed under copper. 

Pure aluminum has a high conductivity for electricity, about 
two-thirds that of copper for conductors of equal cross-section, 
but since the density of aluminum is so much smaller than that 
of copper, an aluminum conductor will weigh less than a copper 
conductor of equal carrying capacity. 

Chemical Properties. Metallic aluminum is really a very 
active element and passes into its compounds with a great 
decrease in free energy, and yet it acts in the main like a rather 
inert metal. It seems not to be acted upon by air or water or 
many other chemical agencies. But this indifference is apparent 
rather than real and is due to the formation of a very thin, 
closely adhering film of the oxide which acts as a sort of varnish, 
and in a great measure protects the metal from further attack. 
Aluminum may be amalgamated by contact with solutions of 
meruric chloride. The aluminum oxide seems to be unable to^ 
clingp' to the surface of the fluid or semi-fluid aluminum amalgam 
and the latter shows the real activity of the metal. It is rapidly 
acted upon by air and water giving hydrogen and aluminum 

Aluminum dissolves in solutions of sodium hydroxide with the 


evolution of hydrogen. The explanation for this action is found 
in the fact that sodium hydroxide will dissolve aluminum oxide, 
and therefore it is supposed that it removes the protecting film 
from the surface of the metal and allows the latter to display its 
activity. Because of this protecting film, massive aluminum is 
but superficially attacked even at high temperatures by the air. 
The powdered metal, however, will burn with a very bright light 
when blown into a flame especially if carried by a stream of oxy- 
gen. It is sometimes used in flashlight powders. Aluminum dis- 
solves readily in hydrochloric acid, not so easily in dilute sulf uric, 
and very slowly in dilute nitric. In the two former cases, hydro- 
gen 'is evolved while in the latter, ammonium salts are produced. 
The difference in the rate of attack is probably connected with 
the film mentioned above. Another phenomenon which is 
probably due to this is, that if a piece of aluminum is made anode 
in certain salt solutions, sodium sulfate, carbonate, phosphate, 
etc., it will not allow a current to pass until a certain potential is 
reached, which in the case of the phosphate is 300-400 volts; 
while if it is made the cathode, the current passes easily at a low 
voltage. The aluminum electrode changes its properties very 
rapidly with the change in the direction of the current, and this 
property may be made the basis of a method for changing an 
alternating into a direct current. It is also of use in certain kinds 
of lightning arresters. 

Aluminum is trivalent in all of its compounds. Its hydroxide 
acts both as an acid and as a base and therefore cannot be strong 
in either way, consequently both classes of salts are hydrolyzed 
to a certain degree. 

The atomic weight of aluminum is 27.1. 

Goldschmidt Process. The most striking chemical property 
of aluminum is its great tendency to combine with oxygen. As 
mentioned above this is generally concealed by the film of oxide, 
but under the proper conditions, aluminum will reduce all other 
oxides except magnesium. To bring this about, the oxide to 
be reduced is mixed in the proper proportions with powdered 
aluminum (Fig. 66) and the mixture heated at one point to a 
high temperature by burning magnesium or a mixture of mag- 
nesium and potassium chlorate, when a very vigorous reaction 
takes place which spreads through the entire mixture. The 



temperature attained is usually high enough to melt the other 
metal and the aluminum oxide. Naturally it varies with the 
oxide reduced. The process was invented by Goldschmidt and 
called by him " Aluminothermy " and the mixtures "thermite." 
They are used both to get high temperatures and pure metals. 
The equation for the reaction with ferric oxide is, 

Fe 2 3 + 2A1 = 2Fe + A1 2 O 3 

This reaction produces a temperature of 3,000 and is often 
used for welding operations and in the repair of breaks in large 
castings or forgings. It can be used almost anywhere and has 
proved to be of very great benefit for making repairs at a dis- 
tance from a machine shop. With sulfides, a very similar reac- 
tion takes place, aluminum 
sulfide, A1 2 S 3 , and the metal 
being formed. 

Aluminum Oxide and 
Hydroxide. Aluminum 
oxide, A1 2 O 3 , or alumina as it 
is often called, is found in 
nature in the mineral corun- 
dum. When colored red by 
chromium it is called ruby, 
when blue, sapphire. Corun- 
dum mixed with magnetite is 

>...>* Magnesium Ribbon 

and Potassium 

Mixture of Powdered 
'Aluminum ctnd the 
Oxide to be reduced 

FIG. 66. 

called emery. Corundum is 
next to diamond and carborun- 
dum (silicon carbide) in hardness, and is used for making wheels 
for grinding. Aluminum oxide which has been fused in an elec- 
tric furnace is identical with corundum and is now being used 
instead of the mineral under the name "alundum." Artificial 
rubies which are identical with the natural stones, are made by 
fusing a mixture of aluminum oxide and a chromium compound. 
Aluminum hydroxide is found in nature in the mineral hydrar- 
gillite, A1(OH) 3 , and may be made in the laboratory by the 
addition of a solution of a base or of a carbonate to a solution of 
an aluminum salt.' It is not soluble in water, but is dissolved by 
acids and by strong bases, resembling the hydroxides of beryllium 
and zinc in this respect. It differs from zinc hydroxide in that 


it is very slightly soluble in ammonium hydroxide. The hy- 
droxide dissolves in acids because of the usual dissociation of the 
bases into the cation and hydroxylthe concentration of the latter 
being decreased by the hydrogen. The bases dissolve it, because 
it splits off the hydrogen ion and forms the aluminate ion AlOs", 

A1(OH) 3 <=H+ + H 2 + A10 2 - 

The hydroxyl ion of the base unites with the hydrogen ion 
and decreases its concentration. The salts formed are alumi- 
nates. Sodium aluminate and the corresponding potassium salt, 
NaAlC>2 and KA102, are soluble in water, but are highly hydro- 
lyzed because aluminum hydroxide is a very weak acid. Most 
of the aluminates are insoluble and many are found in nature; 
those of the bivalent metals are called spinels from the magnesium 
aluminate, spinel, Mg(AlO2)2, which is the typical mineral of the 

Because of the great weakness of aluminum hydroxide as a 
base, aluminum carbonate does not exist, and aluminum hy- 
droxide is precipitated upon the addition of carbonates to solu- 
tion of aluminum salts. Obviously salts of other weak acids 
will work in the same way and A1(OH) 3 is precipitated by 
solutions of alkali sulfides and cyanides. 

Aluminum Chloride. Aluminum chloride, A1C1 3 , is the most 
important of the artificially prepared halogen compounds of 
aluminum. It may be obtained in the form of a crystalline 
hydrate, A1C1 3 -6H 2 O, from solutions prepared by dissolving the 
hydroxide or metal in hydrochloric acid. When heated to drive 
off the water, the chloride is completely decomposed, aluminum 
oxide and hydrochloric acid being formed. The anhydrous 
chloride may be obtained by heating the metal in chlorine or 
in hydrogen chloride. It is a white crystalline solid which 
sublimes without melting and fumes when exposed to moist air. 
The anhydrous salt is used in organic chemistry as a catalytic agent. 

Because aluminum hydroxide is a weak base the solutions of 
aluminum salts, including the chloride, are acid in reaction from 
hydrolysis, and generally can be kept clear only in the presence 
of an excess of acid. 

Aluminum Sulfate. Aluminum sulfate, AWSOOs'lSH^O, is 
made on a large scale from sulfuric acid and either aluminum 


hydroxide or aluminum silicate, H^A^SiO^a'H^O or kaolin as 
it is often called. At the present time, the hydroxide is more 
commonly used. It is very soluble in water and the solution 
is acid because of hydrolysis. Its trade name is " concentrated 
alum." It is used as a mordant in dyeing, in sizing paper 
to prevent the spreading of ink, and in clarifying water. When 
a solution of potassium sulfate is added to a strong solution of 
aluminum sulfate, octahedral crystals of a double salt named 
alum, KA1(SO 4 ) 2 -12H 2 O, are formed. This is the type of a 
rather large number of double salts which are isomorphous 
and are called alums. To distinguish it from the others, it 
is usually called potassium alum. The alums have the general 
formula M+M +++ (A0 4 ) 2 -12H 2 O in which M + represents any 
member of the alkali metals except lithium and in addition 
ammonium, silver, and monovalent thallium; M +++ stands 
for aluminum, gallium, indium, trivalent iron, chromium, and 
manganese; and finally A may be either sulfur or selenium. 
The aluminum alums are distinguished by prefixing the name of 
the monovalent metal while in the case of the other alums both 
the mono- and trivalent metals are named, for example, potassium 
chrome alum, KCr(SO 4 ) 2 -12H 2 0. The 12H 2 O is an essential 
part of an alum. Double sulfates are known which differ from 
alums only in this respect and yet they are not alums. Since 
the alums are isomorphous they can all form mixed crystals and 
a single crystal may contain a large number of elements. 

Potassium alum is the most important alum ; it was formerly 
the most important aluminum salt but has now been largely 
replaced by the sulfate. It may be prepared by roasting alunite, 
a basic alum found in Italy, Hungary, and in southwestern United 
States and extracting with hot water. The crystals melt or dis- 
solve in their own water at 92. 

Solutions of alum are acid and will react with carbonates and 
bicarbonates with the evolution of carbon dioxide, the precipita- 
tion of aluminum hydroxide and the formation of a solution of 
the sulfates of potassium and the metal of the carbonate. It is 
upon this property that its use in cheap baking powders depends. 
The soluble sulfates formed impart a strong bitter taste to the 
food and the use of such powders is objectionable on this account, 
if for no other. A solution of alum will dissolve a considerable 


quantity of aluminum hydroxide forming the so-called "neutral 
alum," KA1 2 (OH) 3 (S0 4 )2. It is used as a mordant in dyeing 
and as a clarifying agent in the purification of water. The 
aluminum hydroxide which it readily forms is the active agent 
in each case. In the clarification of water the precipitated alu- 
minum hydroxide- attaches itself to the silt suspended in the water 
and also to most of the bacteria present and carries the whole to 
the bottom. 

Alum loses all of its water at 100. It is then known as 
burnt alum, and is used in medicine. 

Mordants. Some dyes will combine directly or enter into 
solid solutions with the fiber of the cloth while others may be 
precipitated as insoluble substances within the fiber, but many 
others do not have either of these characteristics. Such dyes 
require a mordant, i.e., a substance which will attach itself 
firmly both to the fiber and to the dye. Aluminum hydroxide 
is such a substance. The cloth to be dyed is first treated with 
a solution of an aluminum salt in such a condition that it con- 
tains a considerable amount of aluminum hydroxide; neutral 
alum, aluminum sulfate, acetate, or sodium aluminate serve 
well; and it is then boiled with the dye. Dyes more often 
require mordants with cotton than with wool or silk. Aluminum 
hydroxide, even in the absence of the fiber, will take up many 
dyes forming what are known as " lakes." 

Aluminum Acetate. Aluminum acetate, A1(C 2 H 3 O2)3, may 
be made by the action of the sulfate in solution upon lead or 
barium acetate, lead or barium sulfate being precipitated. 
Since it is a salt of a weak acid and of a weak base it is strongly 
hydrolyzed, and on this account is used in mordanting cloth 
and also in making it waterproof. The same result may be 
secured at a lower cost by using aluminum sulfate and sodium 

Aluminum Sulfide. Aluminum sulfide, A1 2 S 3 , cannot be pre- 
pared in a wet way. It is formed by the action of finely divided 
aluminum on sulfides at a high temperature by practically the 
methods of the Goldschmidt process. 

Aluminum Silicate Clay. The rocks which originally com- 
posed the earth's crust were very largely complex silicates and 
almost invariably contained aluminum as one of their constitu- 


ents. By the action of the carbon dioxide of the air and of water, 
these rocks gradually decompose or weather as it is called, and 
aluminum silicate is one of the products. The aluminum silicate 
being practically insoluble and in an extremely fine state of divi- 
sion is carried by flowing water and deposited when the latter 
reaches a quiet lake or the ocean, forming the beds of clay. 
These evidently have a very good chance to become contami- 
nated and are rarely pure. The purest clay, kaolin or china 
clay, H 2 Al2(Si04)2-H 2 O, has probably been formed by the decom- 
position of feldspar, KAlSi 3 8 , the potassium carbonate being 
carried away and the kaolin left where it was formed. Common 
clay contains some calcium carbonate, iron oxide, quartz, etc. 
When comparatively pure it is known as potters earth. Marl 
contains large quantities of calcium carbonate. Ocher, umber 
and sienna are clays colored by oxides of iron and manganese. 
They are used to some extent as pigments. 

Clay, as is well known, becomes plastic when moist and may 
be easily worked into almost any shape desired. When dried and 
then heated to a high temperature it shrinks, becomes quite hard, 
and loses its power of becoming plastic with water. It also 
becomes resistant to chemical change and acquires considerable 
mechanical strength. Because of these properties it is used for 
making brick, pottery, and porcelain. The chemical change 
which takes place during the heating consists in the driving off 
of all of the hydrogen and part of the oxygen in the form of water. 
Pure kaolin cannot be melted in the furnaces for firing this mate- 
rial, but the presence of compounds of the alkalies, calcium or 
magnesium carbonates or especially of oxide of iron causes the 
mixture to fuse partially and become more or less pasty while in 
the furnace, and stronger after the firing. The iron" imparts the 
red color commonly associated with brick. Fire brick is made 
from nearly pure clay and has a high melting-point to which it 
owes its name. Articles made from clay are more or less porous 
after firing, and for many purposes they must be glazed. This is 
accomplished in various ways as will develop. 

Porcelain and Pottery.- Porcelain is the name applied to the 
highest grade of ware made from clay. It is made by grinding 
together very pure kaolin and feldspar in the proper proportions, 
and making up into a plastic mass from which the desired articles 



are formed. After drying, these are fired at such a temperature 
that the feldspar melts and binds the burned kaolin together 
into a hard translucent mass which is nearly homogeneous 
but does not have a glaze. To secure the latter the articles 
are dipped into a suspension of very finely ground feldspar in 
water, the ware becomes coated with a thin layer of the feldspar, 
and after drying is fired once more at a high temperature and is 
then allowed to cool slowly (annealed). The ware is then trans- 
lucent and has a smooth glaze which is practically a glass. It is 
very difficultly fusible and decidedly resistant to chemical action, 
but is attacked by aqueous and fused alkalies. 

Semiporcelain is made from a white plastic clay to which is 
t added ground quartz, flint, or feldspar. After the first burning 
it is porous and is glazed with a mixture of borax, quartz, sodium 
carbonate, and oxide of lead. The ware is usually much thicker 
than porcelain and is not very suitable for chemical purposes. 

Earthen-ware is made from colored plastic clay, the impurities 
of which serve to frit or bind the kaolin together. It is glazed by 
throwing common salt into the kiln toward the close of the burn- 
ing. At the high temperature, the salt volatilizes and reacts 
with the clay and the water vapor present so as to form an easily 
fusible silicate which fills the pores of the ware. Unglazed 
earthenware is largely used in the form of tile, bricks, flower 
pots, etc. 

Ultramarine. A double silicate of sodium and aluminum 
containing some sulfur is found in nature as the mineral lapis 
lazuli which has approximately the composition (NaAlSiO 4 )4- 
Na 2 S2. It has a very beautiful blue color and is used for orna- 
mental purposes. When ground it forms the pigment known as 
ultramarine. This same pigment may be made very cheaply 
by heating together sodium sulfate, carbon, and clay. The 
resulting, blue compound is ground and washed with water. 
Since it is very cheap and is stable toward light, air and alkalies 
it is much used as a pigment in water color paints, wall paper, 
laundry blue, etc. It was formerly used to disguise the faint 
yellow color of granulated sugar. It loses its color on contact 
with even very weak acids and the loss of color is accompanied 
by the evolution of hydrogen sulfide. The blue color of the 
compound is something of a puzzle because its components might. 


.naturally be expected to yield a colorless substance. By varying 
the methods of manufacture, green, violet, and red varieties may 
be obtained. 

Hydraulic Mortars and Cements. When a limestone contain- 
ing something like 8 to 18 per cent, of clay is " burned" in the 
ordinary way the product will slake and form a mortar which will 
set and harden under water and hence is called hydraulic cement. 
A much better product is made by burning a limestone which 
contains 20 per cent, or more of clay until the carbon dioxide is 
driven off. The substance so produced will not slake; but if 
finely ground and mixed with water, will set and harden out of 
contact with the air. This is known as natural cement. 

Portland cement is the best hydraulic cement. It is the finely 
pulverized product made by heating to incipient fusion a properly 
proportioned mixture of silica, aluminum silicate, and calcium 
carbonate to which an addition of not more than 3 per cent, of gypsum 
has been made after the heating. The essential constituents of 
Portland cement are lime, silica, and alumina. Under commer- 
cial conditions some of the alumina is always replaced by ferric 
oxide, and a little of the lime by magnesia. 

Portland cement may be made from almost any materials 
which give the proper mixture of calcium carbonate, silica, and 
aluminum silicate provided they do not contain more than 4 per 
cent, of magnesium carbonate and that the silica is in a very fine 
state of subdivision. The most widely used materials for the 
preparation of this cement are limestone and clay or shale ; these 
are very finely ground and thoroughly mixed in such proportions 
that the mixture shall contain close to 75 per cent, of calcium 
carbonate and 20 per cent, silica, alumina and ferric oxide taken 
together, the remaining 5 per cent, being magnesium carbonate, 
alkalies, etc. Some limestones have approximately this com- 
position, and are called " cement rock' 7 since they require merely 
the addition of a little of a purer limestone or of a suitable clay 
to correct their composition. 

After the components of the cement have been ground very 
fine and thoroughly mixed in the proper proportions, they are 
then fed into the upper end of a long, slightly inclined rotary kiln 
through which they gradually work their way against the flames 
and hot gases of a fire of powdered coal which is blown into the 



lower end of the kiln with a blast of air, and which burns much as 
a gas would do (Fig. 67). 

These kilns are 60 to 150 ft. in length and 6 to 9 ft. in diameter. 
Sometimes they are fired with crude oil, or natural gas, but coal 
is the usual fuel. At the high temperature reached (1,500- 
1,600) the calcium carbonate is decomposed into lime and 
carbon dioxide and the lime reacts with the silica and aluminum 
silicate to form calcium silicates and calcium aluminates which 
partially fuse. The product is termed clinker and comes out 
in hard lumps. It is then ground once more and mixed with 

Ql-2% Gypsum 


: *; Air and Coal 

FIG. 67. 

2 to 3 per cent, of calcium sulfate in the form of gypsum or 
plaster of Paris. This acts as a "retarder" and increases the 
time of setting. When the cement is mixed with water it soon 
sets and then gradually hardens and increases in strength for 
months, and even in some cases for years. The calcium silicates 
and aluminates undergo hydrolysis with the formation of some 
calcium hydroxide and new forms of calcium aluminate and 
silicates. The crystals of these interlace and cause the setting. 
The hardening is produced by the gradual formation of a col- 
loidal mass of monocalcium silicate which cements the crystals 


A mixture in the proper porportions of Portland cement, sand, 
crushed stone and water will harden to a stone-like substance 
called concrete, the usefulness of which is well known to every 
one. In 1914 the United States produced 88,230,000 barrels of 
Portland cement. 


Gallium has an atonic weight of 69.9 and is the third member 
of the aluminum family. Its existence and some of its prop- 
erties were predicted by Mendelejeff. The element is widely 
distributed but always in very small quantities. It is found in 
many samples of zinc blende, in some iron ores, and always in 
bauxite. The metal melts at 30.2, has a bluish-white color and 
a density of 5.9. Toward air, water, acids, and bases it acts as 
aluminum does. It is mostly trivalent in the compounds, 
although a dichloride, GaCl 2 , is known. The hydroxide is both a 
base and an acid. The sulfate forms alums with the sulfates of 
the alkali metals. It is most easily detected by its spark- 


Indium is the fourth member of the group and has an atomic 
weight of 114.8. It occurs in small quantities, as the sulfide in 
certain zinc blends. The metal has a density of 7.1 and is white 
and easily malleable; it melts at 155. It retains its luster in 
air and water. It is dissolved by dilute hydrochloric and sul- 
furic acids, and more readily by nitric acid. It is mostly trivalent 
in its compounds, but a mono- and dichloride are known. 
The hydroxide is soluble both in acids and in the strong bases. 
The sulfate forms alums. The spectrum contains a strong indigo 
blue line and a violet line. 


Thallium is found in many varieties of iron and copper pyrites. 
It also occurs as an essential constituent of a few rare minerals. 
It is somewhat more abundant than gallium and indium. Its 
atomic weight is 204.0. In a general way the metal resembles 
lead, except that it is much more readily oxidized. The density 
of the metal is 11.8. It melts at 301, and boils at 1,515. The 


metal dissolves readily in dilute and concentrated sulfuric acids, 
and less readily in hydrochloric. Thallium forms two series of 
compounds. In one, the thallic series, it is trivalent and re- 
sembles aluminum to a certain degree, but the hydroxide is not 
soluble in bases and the sulfate does not form alums. 

In the other the monovalent or thallous series it presents a 
curious set of relationships. The hydroxide, ThOH, is soluble 
in water and is a strong base. The carbonate, phosphate, sul- 
fate and oxalate are soluble. These properties are very remark- 
able for a heavy metal, and are like those of an alkali metal. This 
resemblance is further strengthened by the facts that the sulfate 
forms the compound .T^SO^MgSO^GH^O, isomorphous with 
the corresponding potassium salt, and also forms alums with the 
ferric, chromic, or aluminum sulfates, which are isomorphous 
with ordinary alum. The fluoride is easily soluble while the 
rest of the halides are but slightly soluble, the solubility decreas- 
ing from the chloride to the iodide. The chloride changes from 
white to violet in the light. The chromate is slightly soluble. 
In all these ways it is like silver. It resembles lead in that the 
chloride is not dissolved by ammonia but is by hot water. 


The members of Group IV, like those of the preceding groups, 
divide themselves into two sub-groups, which may be named 
the titanium and the germanium sub-groups. There is not as 
marked a difference in the properties of the two as was the case 
with the preceding groups, and yet the same general relationship 
holds here. The basic character of the elements increases with 
the atomic weight. The members of the titanium groups, 
which occupy the left-hand column in the table as it is arranged, 
are, on the whole, stronger bases than those on the right-hand 
side, the germanium group. This, it will be recalled, has been 
the case in each of the prreceding groups of metals. The gradual 
weakening in the basic properties, which has been noticeable in 
passing to each successive group, is continued with this and there 
is no really strong base formed by a member of the group, in 
fact, the first two members, carbon and silicon, are distinctly 
non-metallic, and have been discussed among such elements. 
The maximum valence of the members of the group is four. 


The members of this sub-group are titanium, atomic weight, 
48.1; zirconium, atomic weight, 90.6; cerium, atomic weight, 
140.25; and thorium, atomic weight, 232 .4. 


Titanium is a widely distributed and comparatively abundant 
element being present in almost all igneous rocks and forming 
a larger per cent, of the lithosphere than carbon. It does not 
occur free in nature, and its compounds are not collected in great 



beds like the coal beds, so it is not as available as carbon. It 
occurs in a number of minerals, such as rutile, TiO 2 and titanic 
iron ore or ilmenite, FeTiOs. The metal dissolves in acids with 
the evolution of hydrogen, and decomposes steam at a high 
temperature. It forms three series of compounds in which it is 
di-, tri-, and tetravalent; the last are the most stable. Titanic 
hydroxide is rather more of an acid than of a base but acts in 
both ways. Metallic titanium has a great tendency to combine 
with nitrogen, oxygen and sulfur. An alloy of metallic titanium 
with iron, called ferrotitanium, is often used in the manufacture 
of steel. The titanium increases the tensil strength and toughens 
the steel. The carbide is used in special kinds of flaming arc 
lamps, burning titanium carbide as cathode below a copper 


Zirconium is a rare metal, which is found chiefly as zircon, the 
silicate ZrSiO 4 . Its hydroxide is a weaker acid and a stronger 
base than titanium hydroxide. 


Cerium is a rare metal which is chiefly found in cerite, a hy- 
drated silicate, which contains calcium, cerium, and lanthanum, 
and the other rare earths which so closely resemble lanthanum. 
The metal burns more readily than magnesium and the small 
particles thrown off, when it is scratched with a hard object, 
take fire. For this reason an alloy of cerium and iron which 
gives brilliant sparks when rubbed with a file is used as a gas 
lighter. Cerium dioxide is used in connection with thorium 
dioxide in the preparation of incandescent mantles, as will soon 
be described. Cerium forms two series of compounds, in one it is 
trivalent and strongly resembles the rare earths, while in the 
other it is tetravalent, and is like the other titanium metals. 


Thorium is obtained chiefly from monazite, a phosphate of 
cerium and thorium, Thorium hydroxide is the strongest base 


of the titanium series. The nitrate, Th(N0 3 )4*6H 2 is the most 
important salt. It is used in the formation of Welsbach incan- 
descent gas mantles. These are made by dipping a mantle 
woven from ramie fiber into a solution of thorium nitrate along 
with about 1 per cent, of cerium nitrate, Ce(NO 3 ) 4 . After 
drying, the ramie is burned away, and the nitrates are converted 
into the oxides which retain the form of the mantle. A mixture 
of the two oxides in the proportions given is very efficient as a 
light producer, while pure thorium dioxide, or a mixture contain- 
ing more cerium, is not nearly so good. 

Thorium and its compounds are radioactive, but nowhere near 
as highly so as the radium compounds. A discussion of this 
phase of the properties of thorium will be postponed until later, 
when all the radioactive elements may be treated together. 


The members of this sub-group are both di- and tetravalent. 
In their divalent compounds, they are more strongly basic than 
in the other series. 


Germanium, atomic weight, 72.5, is a rare element. Its prop- 
erties agree closely with those predicted by Mendelejeff for eka- 
silicon, germanium being then undiscovered. The element itself 
is distinctly metallic, but its compounds are much more like those 
of the non-metals than the metals. It bears a strong resemblance 
to carbon and silicon, as may be seen from the formulas of a few 
of its compounds, GeH 4 , GeF 4 , GeCl 4 , GeCl 2 , GeHCl 3 (ger- 
manium chloroform), Ge(>2, GeO, GeS2, GeS. The dioxide has 
acid properties, but also dissolves in acids. The sulfide, like 
zinc sulfide, is white. 


The discovery of tin is prehistoric, the metal having been used 
by early man in making his bronze tools and weapons. 

Occurrence and Metallurgy. The chief ore is tin stone, or 
cassiterite, Sn02. This is found in many places, but chiefly in 
the East Indies, Bolivia and Cornwall. It is easily reduced to 


the metallic state by carbon, the greatest difficulty in the metal- 
lurgy of tin being the purification of the oxide before its reduc- 
tion. This is done by crushing the ore and mechanically separat- 
ing the granite or other rocky gangue. The concentrated ore is 
then repeatedly roasted and washed to remove the sulfides 
of arsenic, iron and copper. The purified stannic oxide is then 
reduced in reverberatory furnaces with carbon in the form of 
anthracite coal. The tin so obtained contains some iron and 
arsenic. These combine with a little tin to form a compound 
having a higher melting-point than the pure tin, so that the latter 
may be purified by carefully heating the ingots to a very little 
above the melting-point of tin, and allowing the purer metal to 
run away from the iron arsenic alloy. This process is called 
liquation. The purest tin is that from Banca. The world's 
production of tin in 1914 was 120,400 tons. 

Physical Properties. Tin is a silvery white highly lustrous 
crystalline metal, which exists in several enantiotropic modifica- 
tions. The one which is the most familiar is stable from 18 to 
161 and is known as tetragonal, or ordinary tin. Below 18 it 
becomes meta-stable, and may pass over into a gray powdery 
modification, with a considerable increase in volume, since the 
density of the gray tin is 5.8, while that of the ordinary tin is 
7.3. This is, of course, accompanied by the injury or destruction 
of any object made of the tin. The change is known as the 
"tin disease" and has caused the gradual decay of organ pipes, 
etc., in the northern part of Europe where the temperature 
averages lower than 18. This change is most rapid at 50. 
The transformation is very easily suspended and in the absence 
of the gray modification, ordinary tin will remain at tempera- 
tures somewhat below 18, practically indefinitely. At 161, 
tetragonal or ordinary tin changes into rhombic, which is stable 
up to the melting-point, 232. Tin is very malleable and may be 
beaten or rolled into thin sheets, known as tin foil. Cast tin is 
crystalline and when bent gives out a peculiar sound known as 
the "cry of tin." Tin plate or "tin," as it is generally called, is 
made by dipping carefully cleaned sheets of mild steel into 
molten tin. The tin forms an alloy with the iron and this, in 
turn, becomes covered with a thin layer of practically pure tin. 
Such plate owes its usefulness to the fact that metallic tin is 


unaffected by air, water, or by weak acids in the absence of air. 
As is well known, very large quantities of tin plate are used in 
making cans and utensils. The greater part of these soon find 
their way to the rubbish heap and the tin is lost. This waste is 
regretable because the supply of tin ore seems to be limited. 
Attempts are being made with somewhat encouraging success to 
recover the metal from such waste plate. One very serious 
drawback is the cost of collection. 

There are in use two important methods for the recovery of 
tin. In one the tin plate is made anode in a solution of sodium 
hydroxide. The tin dissolves as sodium stannate and is deposited 
on the cathode. In the other, the plate is exposed to the action 
of chlorine which attacks the tin much more easily than the iron. 
Each of these methods works well on the clean scrap left in the 
manufacture of tinned articles and about one-fourth of the tin 
used in this country is recovered in these ways. 

A piece of tin plate which has once got a hole scratched 
through the tin down to the iron will rust more rapidly at that 
place than a piece of pure iron. This is due to the fact that in 
the presence of water an electric battery is formed with the iron 
acting like the zinc in an ordinary cell, with the result that it is 
rapidly corroded. 

Tin is a component of many alloys; bronze has been given; soft 
solder contains 50 per cent, each of lead and tin; pewter, 25 per 
cent, lead; Britannia metal 10 per cent, antimony and some 
copper. Tin amalgam was formerly used to back mirrors. 

Chemical Properties. Tin forms two series of compounds, the 
stannous, in which it is divalent, and the stannic, in which it is 
tetravalent. Stannous hydroxide is more basic than stannic 
and its salts are less hydrolyzedln water solutions than the stan- 
nic salts, but each hydroxide is soluble both in bases and in acids 
and therefore each is a weak base. Neither is a strong enough 
base to form a carbonate. The stannous salts are good reducing 
agents, passing readily into the stannic compounds. 

The metal is not altered by air or water at ordinary tempera- 
tures, but oxidizes easily at higher temperatures. It dissolves 
readily in hydrochloric acid with the evolution of hydrogen and 
the formation of stannous chloride, SnC^. Hot concentrated 
sulfuric acid forms stannous sulfate, SnS0 4 , and sulfur dioxide. 


Cold dilute nitric acid slowly acts upon it forming stannous 
nitrate, Sn(NO 3 )2, and ammonium nitrate, NH 4 N0 3 , while con- 
centrated nitric acid acts rapidly, forming white insoluble meta 
stannic acid, H 2 Sn03, and nitrogen peroxide, 

Sn + 4HN0 3 = H 2 SnO 3 + 4N0 2 + H 2 O 

Tin will dissolve in sodium or potassium hydroxide, giving 
hydrogen and a stannate, such as K 2 Sn0 3 . 
The atomic weight of tin is 118.7. 


Stannous Oxide, Hydroxide and the Stannites. Stannous 
oxide, SnO, is a black powder, which will burn in the air, forming 
stannic oxide. It is made by heating stannous oxalate, SnC 2 04, 
which decomposes into the oxide, carbon monoxide and carbon 
dioxide. The corresponding hydroxide is precipitated from a 
solution of stannous chloride upon the addition of sodium 
hydroxide or carbonate. It is a white, gelatinous precipitate, 
which is soluble, both in acids and strong bases, forming in the 
latter case stannites, such as sodium stannite, Na 2 SnO 2 . These 
solutions are powerful reducing agents and will reduce the com- 
pounds of many of the metals to the metallic state. When a 
stannite solution is boiled a part of the tin is reduced to the metal, 
while the remainder is oxidized to the stannate, as shown below 
for sodium stannite. 

2Na 2 SnO 2 + H 2 = Sn + Na 2 SnO 3 + 2NaOH 

Stannous Chloride. Stannous chloride, or tin salt, SnCl 2 2H 2 O, 
is the most important stannous salt. It is made by dissolving 
tin in hydrochloric acid. Like the other stannous salts, it is 
hydrolyzed and the presence of an excess of acid is necessary to 
prevent the formation of a basic salt, Sn(OH)Cl. Solutions of 
stannous chloride are used by the dyers as mordants, and also 
as reducing agents, in the manufacture of dyes, etc. The use 
of it as a mordant depends upon the hydrolysis of the solution 
and the formation of the hydroxide, which has similar properties 
to aluminum hydroxide in its relation toward dyes. The ac- 
tion of the chloride as a reducing agent is due to the tendency 


of the stannous compound to pass into the stannic. Stannous 
chloride will reduce the compounds of mercury and of the noble 
metals to the metallic state, ferric and cupric salts to ferrous and 
cuprous salts, it will also be oxidized by the oxygen of the 
air; and it is necessary to keep some metallic tin in a stannous 
chloride solution to reduce any stannic chloride which may be 
formed in this way. 

Stannous sulfate and nitrate are less stable than the chloride, 
and are of no practical importance. 

Stannous Sulfide. Stannous sulfide, SnS, is analytically 
important. It is formed as a dark brown precipitate when 
hydrogen sulfide is passed into a moderately acid solution of a 
stannous salt. It is soluble in concentrated hydrochloric acid 
and is reprecipitated upon diluting the solution. It is not 
soluble in solutions of the alkali sulfide unless they contain 
polysulfides which act by first oxidizing the_ stannous jsulfide to 
stannic sulfide, which then dissolves, forming thiostannates, such 
as sodium thiostannate, Na 2 SnS 3 . This formation of soluble 
thio salts relates tin to gold, arsenic and antimony, as will soon 
be seen. 


Stannic Oxide and Hydroxide. Stannic oxide, SnCh, as has 
been mentioned, occurs in nature and is the principal ore of tin. 
It is white when pure, but the natural compound is usually 
colored from impurities. It may be made by heating meta 
stannic acid which decomposes into water and stannic oxide. 
When prepared at a low temperature, it is easily soluble in acids; 
but when ignited, it becomes very difficult to attack. 

Two stannic hydroxides are known, each of which behaves as 
an acid. When dried they both pass into the dioxide. To dis- 
tinguish them, one is called alpha and the other beta stannic 
acid. Alpha stannic acid is formed by precipitating stannic 
chloride solution with ammonium hydroxide, or by carefully 
acidulating an alpha stannate. It is a white gelatinous precipi- 
tate, which gives to water a faintly acid reaction, and which 
dissolves readily in dilute mineral acids and in solutions of sodium 
or potassium hydroxides, forming alpha stannates, Na 2 Sn0 3 -3H 2 O, 
for example. The sodium salt is much used as a mordant in 


calico printing and is called " preparing salt." It is also used to 
render cotton goods permanently fire proof. The cotton is first 
soaked in a solution of alpha sodium stannate, dried, and then 
placed in a solution of ammonium sulfate. In the presence of 
the cation of the weak base, ammonium hydroxide, the stannate 
is hydrolyzed and stannic acid precipitated in the cloth. 

Beta or meta stannic acid is made by acting upon tin with 
concentrated nitric acid. It differs from the alpha acid in that 
it is entirely insoluble in nitric acid, that it swells up but does not 
dissolve in sulfuric acid, and when treated with hydrochloric 
acid forms a compound which is not soluble in hydrochloric acid, 
but which dissolves after the hydrochloric acid is removed and 
replaced by water. The solution so obtained is not like that of 
stannic chloride and gelatinizes when boiled. Sodium hydroxide 
dissolves beta stannic acid and yields a salt, sodium beta stannate, 
Na 2 Sn 5 On4H20; by acidulating this solution, beta stannic acid 
is precipitated. When beta stannic acid is fused with sodium 
hydroxide, the alpha stannate is forrned. 

When solutions of alpha stannic acid in hydrochloric or hydro- 
bromic acid stand for a long time, beta stannic acid is gradually 

Stannic Chloride. Stannic chloride, SnCl 4 , is formed by the 
action of chlorine upon stannous chloride or upon the metal. 
Use is made of this in recovering tin from scraps of the plate by 
the chlorine process. It is a colorless liquid, which boils at 114. 
It fumes in contact with the air. With water a number of 
hydrates are formed. The pentahydrate, SnCl 4 -5H 2 O, is much 
used in dyeing as a mordanting agent, under the name of "oxy- 
muriate of tin." It combines with hydrogen chloride to form 
chlorstannic acid, H^SnCle, several salts of which are known. 
The ammonium salt, (NH^SnCle, was formerly much used as a 
mordant in dyeing under the name of "pink salt," not because 
it had a pink color but because it was used in dyeing pinks. 

Solutions of the chloride in water are very slowly hydrolyzed. 
The stannic hydroxide remains in solution in the colloidal state, 
and is gradually transformed into beta stannic acid. 
, Stannic Sulfide. Stannic sulfide is precipitated as a yellow 
amorphous substance by passing hydrogen sulfide through a 
moderately acid solution of a stannic salt. It is soluble in con- 


centrated hydrochloric acid, and is reprecipitated by diluting 
the solution. It is also soluble in solutions of ammonium sulfide 
or of the alkali sulfides, forming thiostannates, 

SnS 2 + (NH 4 ) 2 S = (NH 4 ) 2 SnS 3 

Upon acidulation with dilute acids, thiostannic acid is formed, 
which at once decomposes to give stannic sulfide and hydrogen 

Analytical Properties. The properties of stannous sulfide and 
those of stannic sulfide, which have just been given, put it in an 
analytical group with arsenic and antimony, which is dis- 
tinguished from that containing mercury, copper, cadmium, bis- 
muth, and lead by the fact that the sulfides of the last mentioned 
metals are not soluble in ammonium sulfide (copper is slightly 
soluble) . In solution, both the stannic and stannous compounds 
are colorless. The stannous are distinguished from the stannic 
by the dark color of the sulfide and by the fact that the former 
are strong reducing agents. 


Lead is the member of the germanium sub-group with the 
highest atomic weight, 207.2. Although it occurs free in nature 
in insignificant quantities, it is so easily smelted from its ores 
that it was well known to the ancients. It was used by the 
Romans for water pipes, a use to which it is still put. 

The chemical relationships of lead are somewhat varied, as is 
often the case with the last member of a sub-group. It is both 
di- and tetravalent. In the divalent state, lead resembles 
barium in several ways, in others it is like silver and thallium. 
The hydroxide, Pb(OH) 2 , acts both as a base and an acid in that 
it dissolves and forms salts with both acids and bases. In this 
way lead is like zinc. The hydroxide is a stronger base than acid, 
so that its salts with the stronger acids are fairly stable in solution. 
The tetravalency of lead shows its relationship to the other 
members of this group, but the resemblance is not very strong. 
The hydroxide, Pb(OH) 4 , is both basic and acidic; but is a very 
weak base, so that its salts, even with the strongest acid, are 
practically completely hydrolyzed. The tetravalent lead com- 
pounds tend to pass to the divalent state and are good oxidizing 




agents, so the oxidation relationships are the reverse of those 
between stannous and stannic tin. 

Occurrence. The principal ore of lead is galena, lead sulfide, 
PbS, which contains, when pure, 86.57 per cent, of lead. It is 
found in nearly all countries and very often carries silver, either 
as finely divided native silver, or as isomorphous sulfide, dis- 
seminated through it. 

Other common ores of lead are the carbonate, cerusite, PbCO 3 , 
and sulfate, anglesite, PbSO 4 . They are formed as decomposi- 
tion or oxidation products from galena, and are consequently 

found near the surface in mines 
which lower down yield galena 
almost exclusively. 

Qm :||a Metallurgy. Since lead sulfide is 

I \fl flT I ^ e chief lead ore, the extraction of 

I I ||fl lead from its ores is largely devoted 

to the treatment of this compound. 
This is accomplished entirely by 
smelting or dry processes, and the 
blast-furnace process is used almost 

The lead blast-furnace, Fig. 68, 
is rectangular in shape and is from 
15 to 25 ft. high, 10 to 20 ft. long 
and from 3 to 5 ft. wide at the 
bottom. A row of blast pipes, or 
tuyeres, along each side of the 
furnace, serve to introduce the air 
blast, which has a pressure of from 
1 to 3 Ib. 

The lead ore mixed with coke to furnish heat for smelting, 
and sufficient limestones to flux off the silica from the gangue, is 
charged into the furnace. A part of the lead ore is charged raw, 
that is, unroasted, while another portion is previously roasted 
to lead oxide in a separate roasting furnace. This roasted ore 
usually has in it a considerable quantity of iron oxide, which is 
utilized to furnish iron for reducing lead sulfide, if not, iron oxide 
ore is added for this purpose. 

Some of the carbon in the charge reacts with some of the lead 
oxide to form lead and carbon dioxide. 

FIG. 68. 


2PbO + C = 2Pb + CO 2 

Most of the remaining carbon burns to carbon monoxide which 
reduces the rest of the lead oxide. 

PbO + CO = Pb + CO 2 
and also the iron oxide 

Fe 2 O 3 + 3CO = 2Fe + 3CO 2 

The iron formed in this way reacts with the lead sulfide of the 
raw ore to form metallic lead and ferrous sulfide. 

PbS + Fe = Pb + FeS 

Any copper in the ore is dissolved by the ferrous sulfide to form 
a matte which may be smelted for the copper. 

The lead tapped from the blast furnace is far from pure. It is 
called work lead or base bullion and contains besides the gold and 
silver of the ore, considerable quantities of antimony, bismuth, 
copper, and other impurities. Most of the latter can be re- 
moved by melting the lead in large pots and keeping it molten 
for several hours, during which period it is frequently stirred. 
The impurities oxidize more readily than the lead and the 
oxides rise to the top as a scum or dross, which is easily skimmed 
off. Some lead is lost by this process, but the resulting refined 
product is quite soft and malleable. 

This process does not remove gold and silver or all the copper. 
If these are in the lead, it is further treated by the Parke Process 
(see p. 379) to recover the gold and silver. 

Not only does the lead smelting process save the gold and silver 
occurring in the lead ores, but it also allows the smelting of other 
gold and silver ores containing little or no lead, with the lead 
ores, to recover the precious metals and the lead. 

Lead is being successfully purified on a large scale by the 
Betts' electrolytic process. The anodes are made of work lead, 
the cathodes of pure lead and the electrolyte of lead fluosili- 
cate, PbSiF 6 , in solution to which a little glue has been added. 
This latter addition has the effect that the lead is deposited in 
a more coherent form than in its absence. The process pro- 
duces very pure lead and at the same time recovers the precious 



metals in the work lead. The production of lead for 1915 was 
565,000 tons in this country. 

Physical and Chemical Properties. Lead is a bluish-gray 
metal which is very soft and has but little tensile strength; its 
density is 11.4 and its melting-point is 327. At temperatures a 
little below its melting-point, it softens enough so it can be 
forced by hydraulic pressure into the form of pipes which are used 
in plumbing and as a cover for electric cables. There are several 
valuable alloys containing lead; of these, solder, pewter, Britannia 
metal have already been given. Type metal and hard lead 
contain antimony, as does Babbit metal which has tin in addition. 
Shot usually contains a small amount of arsenic. Lead is de- 
cidedly resistant to chemical action and on this account is much 
used in chemical work, as for example in the construction of the 
lead chambers for the manufacture of sulfuric acid. It is but 
slightly acted upon by the oxygen of the air at ordinary tempera- 
tures, and soon becomes covered with a protecting coat of a 
basic carbonate. At higher temperatures it is rather rapidly 
oxidized, forming litharge, PbO, or red lead, Pb 3 04. In the pres- 
ence of air and pure water lead forms the hydroxide, Pb(OH) 2 , 
which is somewhat soluble. On this account lead pipes should 
not be used for carrying drinking water, because all compounds 
of lead are poisonous, and when taken, in repeated small doses, 
gradually accumulate in the system and produce a very serious 
condition known as chronic lead poisoning or " painter's colic." 
If the water contains carbonates or sulfates in solution, the pipes 
become coated with the difficultly soluble carbonate or sulfate of 
lead and the danger of using the water is diminished. However, 
in modern plumbing the use of lead is almost entirely confined to 
the waste pipes. 

Lead stands on the dividing line between the metals which are 
dissolved by the ordinary acids with the evolution of hydrogen 
and those which require an oxidizing agent stronger than hydro- 
gen in the ionic state. It is not dissolved by cold dilute, but is by 
boiling concentrated hydrochloric acid. Sulfuric acid, hot or 
cold, is almost without action until the acid reaches a concentra- 
tion of 80 per cent, or more when the lead is rather rapidly 
attacked, especially by the hot acid. Nitric acid dissolves lead 
readily forming the nitrate, but oxides of nitrogen are evolved 


rather than hydrogen. Acetic acid with the aid of oxygen of the 
air dissolves lead, a fact which is of importance in connection 
with the manufacture of white lead. 

.Oxides and Hydroxides. Lead forms five oxides, having the 
following formulas, Pb 2 0, PbO, Pb 2 O 3 , Pb 3 4 , and PbO 2 . Of 
these the first, the sub-oxide, and the third, the trioxide, are of 
but little importance. The others all serve some very useful 

The monoxide or litharge is prepared by heating lead in the 
air. This is done at one stage in smelting of lead and silver ores; 
the process is called cupellation. Litharge is a crystalline sub- 
stance of a reddish-yellow color. The reddish tinge is due to the 
presence of a little red lead, Pb 3 O 4 . The monoxide is the only 
stable oxide at temperatures above 600, and all the others pass 
into it when heated to this temperature in the air. It is very 
easily reduced to the metal. Litharge is soluble in acids, forming 
salts of divalent lead which although somewhat hydrolyzed are 
in the main fairly stable in solution. It is also soluble in sodiunf 
hydroxide solution, forming sodium plumbite, NaJbQ^ These 
facts show that "the hydroxide, Pb(OH)2, which is undoubtedly 
formed before the oxide is dissolved by either the acids or bases, 
acts both as a base and as an acid. The hydroxide is, however, 
a stronger base than an acid and is strong enough to form a car- 
bonate. The oxide is used in glazing pottery, making glass, and 
in preparing other compounds of lead. 

Lead hydroxide, Pb(OH) 2 , is a white substance slightly soluble 
in water. It is precipitated upon the addition of a solution 
of an alkali to a solution of a lead salt. It is, as indicated above, 
soluble in acids and bases, and yields the same compounds as 
the oxide of lead. 

Red lead or minium, Pb 3 O 4 , is made by carefully heating 
finely powdered lead monoxide or carbonate in the air to a 
temperature not higher than 545. It decomposes at higher 
temperatures into oxygen and monoxide. 

When treated with acids, red lead yields lead dioxide and salts 
of divalent lead, 

Pb 3 Q 4 -f 4HNO 3 = 2Pb(NO 3 ) 2 + PbO 2 + 2H 2 O 
Because of its fine red color, red lead is used as a pigment; it also 


seerns to have a specific action in protecting iron from corrosion. 
It is further used in the manufacture of glass. 

Lead Dioxide. Lead dioxide is the most stable and important 
of the tetravalent compounds of lead. It is a dark brown powder 
which may be formed as described above, but is usually pre- 
pared by the action of bleaching powder upon a solution of lead 

Na 2 Pb0 2 + CaCl 2 + H 2 O = PbO 2 -f 2NaOH + CaCl 2 

The dioxide is insoluble under these conditions and is precipi- 
tated. It has some tendency to act as an acid-forming oxide 
as is shown by the following. When fused with sodium or 
potassium hydroxide, it yields soluble meta-plumbates such as 
K 2 PbO 3 -3H 2 O, analogous to the meta-stannates. Red lead is 
regarded as being lead ortho-plumbate, and on this basis should 
be written Pb 2 PbO 4 . 

Lead dioxide is a strong oxidizing agent. When heated with 
hydrochloric acid, it acts like manganese dioxide and liberates 
chlorine, lead chloride, PbCl 2 , being formed at the same time. 
A few plumbic salts may be made by dissolving the dioxide in 
the cold concentrated acids but they all hydrolyze very easily 
upon dilution and reprecipitate the dioxide. 

From the above it is seen that the dioxide is both basic and 
acidic, but is very weakly basic. 

Lead Chloride. Lead chloride, PbCl 2 , is a white crystalline 
compound which is about three times as soluble in boiling water 
as. in water at room temperature. Advantage is taken of this 
in qualitative analysis in separating lead chloride from silver 
and mercurous chlorides, which are but slightly soluble in either 
hot or cold water. 

Lead Nitrate. Lead nitrate, Pb(N0 3 ) 2 , and acetate are almost 
the only lead salts of the common acids which are freely soluble in 
water. The nitrate is made by dissolving litharge or the metal 
in nitric acid, and crystallizing the product. It forms white 
anhydrous octahedra which are isomorphous with barium 
nitrate. As is to be expected from the weakness of lead hy- 
droxide, a solution of the nitrate is acid in reaction. 

Lead Acetate. Lead acetate, Pb(C 2 H 3 O 2 ) 2 -3H 2 O, or sugar of 
lead, is a white crystalline salt, which is easily soluble in water. 


The solution has a sweet taste, which gives to the salt its common 
name. By boiling a solution of lead acetate with litharge a 
soluble basic salt is formed, Pb(OH)(C 2 H 3 O 2 ). This is alkaline 
in reaction. Both the normal and the basic salts, when in 
solution, yield precipitates of the carbonate when acted upon 
by carbon dioxide; the action is not complete and stops when 
the concentration of the acetic acid has reached a certain point. 
This reaction is of importance in connection with the prepara- 
tion of " white lead." 

Lead Carbonate and White Lead. Normal lead carbonate, 
PbCOs, is found in nature as the mineral cerusite and may be 
prepared in the laboratory, but the artificial product is usually 
a basic salt. Both the normal and basic salts are but very 
slightly soluble in water. 

White lead is a basic carbonate which has approximately the 
composition shown by the following formula, Pb3(OH) 2 (CO 3 ) 2 . 
It is much used as a pigment and is prepared in various ways. 
These nearly all consist in getting the lead in solution with 
acetic acid, and precipitating the carbonate with carbon dioxide, 
thereby recovering the acetic acid to be used to dissolve more lead 
and thus making it act as a catalyzer. 

The white lead which is claimed to be the best is manufactured 
by what is known as the old Dutch process. ,In principle, the 
process consists in the solution of the lead in acetic acid, with 
the aid of the oxygen of the air, and the precipitation of the 
carbonate from this by carbon dioxide, the whole process taking 
place slowly at temperatures only a few degrees above that of 
the room. The equations are, 

2Pb + 4HC 2 H 3 O 2 + 2 = 2Pb(C 2 H 3 O 2 ) 2 + 2H 2 O 
3Pb(C 2 H 3 O 2 ) 2 + 4H 2 O + 2C0 2 = Pb 3 (OH) 2 (C0 3 ) 2 + 6HC 2 H 3 O 2 

The lead to be " corroded" is cast into the form of perforated 
disks called "buckles," Fig. 69, and these are placed in the upper 
part of a glazed earthenware pot, about 8 in. across. In the 
bottom of the pot is some crude acetic acid, and about a third of 
the way up, the sides of the pot are pierced by a number of holes 
to allow circulation of air. A great number of these pots are 
placed side by side upon a layer of spent tan bark. They are 
then covered by boards and another layer of tan-bark, and upon 



the whole another tier of pots is placed, which are covered in 
the same way. This is continued until a pile of many tiers of 
pots has been obtained. Bacteria grow in the tan-bark and 
produce carbon dioxide from the oxidation of its materials, and 
the heat so generated warms the whole to a temperature favor- 
able to the process. The vapors of the acetic acid, the oxygen of 
the air, and the carbon dioxide from the tan-bark being in con- 
tact with the lead bring about the reactions given above. The 
process takes 5 or 6 weeks. At its close the heaps are taken down, 
the lead carbonate is separated from the unchanged lead by 
mechanical means and ground very fine while moist, dried, and 
then ground in linseed oil, making the pasty pigment known as 
white lead. Sometimes the white lead is not thoroughly dried 

FIG. 69. 

before it is ground in the oil; the result is a slightly cheaper paint 
of distinctly inferior quality, which is much more rapidly at- 
tacked by atmospheric agencies than that made from the dry 

White lead is costly, poisonous and hence dangerous to manu- 
facture and use, and is so subject to atmospheric agencies, 
particularly hydrogen sulfide, that it is not suitable for use in 
cities; and yet it has held its own against other white pigments 
which do not have these objections. The reason for this is that 
it has great covering power, i.e., it is very opaque. This is due 
to the fact that the particles of white lead, although they are 
transparent, reflect a good deal of light from each of their 
surfaces, and any given layer of paint contains so many overlying 
particles, each one reflecting a part of the light which penetrates 


those above that practically all of the light is turned back, and 
the paint is decidedly opaque. 

It is a well-established fact that a cheap red barn paint will 
last longer than a high-priced white house paint, when both are 
mixed with equally good oil. The explanation offered is that 
the pigment in the red paint is chemically inert, while the white 
lead in the house paint gradually converts the oil into a more 
or less soluble soap which washes off. The oil used in paints 
must be what is known as a " drying oil" and is usually linseed 
oil. A "drying oil" does not evaporate as water or gasoline 
does, but takes up oxygen from the air and is oxidized to a tough 
elastic substance which binds the pigment firmly to the painted 
surface. "Boiled" linseed oil has been heated with lead oxide 
or manganous borate, these dissolve and catalyze oxidation of the 

Lead Chromate. Lead chromate, PbCr0 4 , is insoluble and is 
formed by bringing together in solution a soluble lead salt and 
a chromate. It has a fine yellow color and is used as a pigment 
and as a dye under the name of chrome yellow. 

Lead Sulfate. Lead sulfate, PbSO 4 , is one of the few sulfates 
which are but slightly soluble. In this respect lead resembles 
barium. Since it is so slightly soluble, it is easily prepared by 
precipitation. It is less soluble in dilute sulfuric acid and in 
alcohol, but more soluble in dilute nitric and in concentrated 
sulfuric acid than in pure water. The decrease in solubility 
upon the addition of dilute sulfuric acid is, of course, due to the 
increase in the concentration of the sulfate ion. The increased 
solubility in concentrated sulfuric acid is probably due to the 
formation of the acid sulfate. The increase in solubility in 
nitric acid is ascribed to the fact that nitric acid is a stronger acid 
than sulfuric, and hence the concentration of the sulfate ion is 

A basic lead sulfate (PbSO^PbO known as sublimed white 
lead is used for a pigment. This compound is made by putting 
a mixture of finely powdered galena, PbS, and carbon on a brisk 
coke fire in a furnace. The lead sulfide is oxidized to the basic 
sulfate which volatilizes at the temperature of the fire, condenses 
in the flue and is caught in bags. It is very fine and white, 
and is chemically much more inert than white lead. 


Lead Sulfide. Lead sulfide, PbS, occurs in nature in cubical 
crystals, which have a gray metallic luster. It is the principal 
ore of lead and is called galena. In the laboratory it may be 
prepared as a black precipitate by passing hydrogen sulfide into 
an acid solution of a lead salt. It is insoluble in cold dilute 
acid, in sodium hydroxide or the alkali sulfides. It is soluble 
in concentrated hydrochloric acid, being reprecipitated upon 
dilution, and also in boiling dilute nitric acid. In the latter case, 
the solution results because of the oxidation of the sulfur ion to 
the element or to the sulfate. 

The Storage Battery. A storage battery is a voltaic cell, 
which, after it has worked for a time and produced electrical 
energy at the expense of the free energy of the reactions within 
the cell, may be restored to its original condition by passing a 
current of electricity through it in the reverse direction. It is 
a reversible cell. One of the most practical forms consists of 
two sets of lead plates, the one coated with lead dioxide and the 
other with spongy lead, the plates being placed in dilute sulfuric 
acid. The oxidizing agent in this case is the tetravalent lead 
of the lead dioxide, which gives up two positive charges to the 
lead electrode, upon which it is placed, and these flow through the 
connecting wire to the other lead electrode and there oxidize 
the finely divided spongy lead, which is the reducing agent, to 
the divalent ion. The divalent lead ion is formed at each elec- 
trode and passes into the nearly insoluble sulfate as fast as it is 
produced. After the cell is discharged, each electrode is coated 
with lead sulfate, instead of lead dioxide, or spongy lead. The 
cell may now be " charged" by passing a current of electricity 
in the reverse direction. At the anode, or dioxide plate, the 
divalent lead ion receives two positive charges and is changed 
to plumbic sulfate, which at once hydrolyzes and gives the 
dioxide. At the cathode, or spongy lead plate, the divalent lead 
ion is reduced to the metal which comes out in a spongy condi- 
tion, and after somewhat more electricity has been passed 
through than was taken out, the cell is restored to its original 
condition. The excess of electricity is due to unavoidable losses. 
The reactions may be represented by the following equations, 
at the positive or dioxide plate on discharge, 


Pb0 2 + 2H 2 SO 4 ^PbSO 4 + 2H 2 O + S0 4 ~ " + 

in which <S> <S> represents the two + charges given up by the 
tetravalent lead and passed through the wire to the negative or 
spongy lead plate, where the following reaction takes place, 

Pb + SO 4 - - + <=*PbSO 4 

On charging, these equations are reversed. 

Analytical Properties of Lead. Divalent lead forms a diffi- 
cultly soluble chloride which puts it in an analytical group with 
silver and mercurous mercury; from these it is distinguished by 
the solubility of lead chloride in hot water. The very slight 
solubility of lead sulfide in dilute acid and in alkali sulfides places 
it in an analytical group with mercuric mercury, copper, cadmium 
and bismuth. From these it is distinguished by the very slight 
solubility of its sulfate, all the other sulfates of the group being 
easily soluble. The properties of the chromate, and hydroxide 
are also of analytical importance. 


Ta<-Nd<- Cb<- V<-N->P-As-Sb->Bi 

Vanadium Sub-group Arsenic Sub-group 

The members of Group V are polyvalent elements with a 
maximum valence of five. The group may be divided into two 
sub-groups, the first consisting of the rare and comparatively 
little known elements, vanadium, columbium, neodymium, and 
tantalum, is called the vanadium group, while the other contain- 
ing the well-known elements nitrogen, phosphorus, arsenic, 
antimony, and bismuth may be called the arsenic group. The 
members of the arsenic sub-groups form a particularly well- 
marked family and change their properties in a very regular 
manner with increasing atomic weight. The tendency ob- 
served with the preceding groups for the basicity to decrease 
from group to group is continued here. Nitrogen and phosphorus 
~are entirely non-metallic in their chemical and physical proper- 
ties; arsenic shows principally non-metallic characteristics with 
some few of the properties of a metal ; antimony is both a base 
and an acid-forming element, while bismuth is almost exclusively 

Every member of Group V forms an acidic pentoxide; the 
lower oxides are more basic, or at least less acidic in their 


From the relations found in the other groups, it might be 
expected that the members of this family would be more basic 
than the corresponding ones of the arsenic group. Vanadium 
meets this expectation for it is distinctly more basic than arsenic ; 
while columbium and tantalum are apparently rather more 
acid than base-forming elements, but their compounds have not 
been thoroughly studied. 


GROUP V 427 

These three metals all combine very vigorously with oxygen 
and have high melting-points. 

Vanadium, atomic weight 51, forms five oxides, V 2 0, VO, V 2 3 , 
V02, V20 5 , analogous to the five oxides of nitrogen. The second 
and third of these are base-forming and yield the vanadous and 
vanadic salts, while the fourth and fifth are acid anhydrides and 
give the vanidites and vanidates respectively. It forms the 
following chlorides: VC1 2 , VC1 8 , VC1 4 , VOC1 2 and VOC1 3 . The 
group VO is very interesting because it acts like a complex di- 
and trivalent cation, forming, for example, vanadyl sulfate, 
VOS0 4 , and di-vanadyl sulfate, (VO) 2 (S0 4 ) 3 . VO++ is called 
vanadyl and V0 + + + di-vanadyl. 

Columbium, Cb, atomic weight 93.5, and tantalum, Ta, atomic 
weight, 181.5, are very rare and hard to separate. Their prin- 
cipal compounds are the columbates and tantalates. 

The atomic weight of neodymium, 144.3, seems to place this 
element in this group but its properties are very much like those 
of cerium with which it is always found associated in nature. 
As yet neither the element nor its compounds have found any 


The members of this family are very similar to nitrogen and 
especially to phosphorus; a proper grasp of their chemistry 
may be most readily obtained by comparison with these elements, 
noting the points of similarity and of difference. At the end of 
this chapter the whole family will be summarized in the form of 
a table. 


Occurrence. Arsenic occurs free in nature and also in com- 
bination, chiefly in the sulfide of the metals in which it replaces a 
part of the sulfur, arsenical pyrite, FeAsS, may be given as an 
example of such a compound. It is also found as the trioxide, 
As 2 O 3 , and as the two sulfides orpiment, As 2 S 3 , and realgar, 
As 2 S 2 . The ores of many metals very often contain arsenic, and 
the removal of this element is one of the serious problems 
connected with their metallurgy. 


Preparation. The element arsenic is obtained on a manu- 
facturing scale by heating arsenical pyrite which decomposes 
into ferrous sulfide and arsenic, 

4FeAsS = 4FeS + As 4 

It may be prepared in the laboratory by reducing the trioxide 
with charcoal. 

As 4 O 6 + 3C = As 4 + 3CO 2 

Physical and Chemical Properties. Arsenic forms at least 
three solid modifications, ordinary or metallic arsenic, yellow 
arsenic, and a black so-called amorphous form. One of these, 
the ordinary or metallic arsenic is very much like red phosphorus, 
another, the yellow arsenic, strongly resembles yellow phos- 
phorus. Ordinary arsenic is steel gray in color, very brittle and 
sublimes at 450 without melting. Like red phosphorus, it is 
not dissolved by carbon disulfide, nor does it phosphoresce in the 
air unless the temperature is raised. When heated, it burns 
forming As 4 O6- It is not attacked by acids with the evolution 
of hydrogen. Nitric acid, however, dissolves it forming oxides 
of nitrogen and arsenic acid, H 3 AsO 4 . | It combines directly with 
the halogens, with sulfur and with many metals. 

Yellow arsenic is made by rapidly cooling the vapor of ordinary 
arsenic. It is crystalline, sulfur-yellow in color and much more 
volatile than the ordinary form. Like yellow phosphorus, 
it is soluble in carbon disulfide and phosphoresces in the air 
at ordinary temperatures. It is even more poisonous than the 
metallic modification. 

The molecular weight of the vapor of arsenic at temperatures 
below 600 indicates a formula of As 4 while at 1,700 it corre- 
sponds to As2. 

A black mirror-like modification of arsenic, which is deposited 
upon cooling the vapors under such conditions that the yellow 
is not formed, is considered to be another modification of arsenic, 
and was formerly held to be amorphous, but it has been shown 
to be crystalline. 

Arsenic forms two series of compounds; in the one it is triva- 
Jent, and inrthe other pentavalent^ There are a few compounds 
in which it is apparently divalent. 

GROUP V 429 


Arsine. Arsenic shows another point of similarity to nitrogen 
and phosphorus in the formation of the hydrogen compound 
arsine, AsH 3 , analogous to ammonia and phosphine. It will be 
recalled that ammonia has a great tendency to combine with 
acids for the formation of ammonium salts, but that in spite of 
this, these salts were practically completely dissociated when in 
the form of vapor. Phosphine showed a much smaller tendency 
to form such salts and only a very few were stable, even in the 
solid state. If this tendency toward instability of the salts 
should continue to increase with the atomic weight, arsine could 
scarcely be expected to form any compounds of this character, 
and as a matter of fact none is known. Ammonia is fairly stable, 
but is largely decomposed at higher temperatures into its ele- 
ments; phosphine is less stable and arsine still less. 

Arsine may be formed by the action of water upon calcium 
arsenide, Ca 3 As 2 , 

~Ca 3 As 2 + 6H 2 O = 2AsH 3 + 3Ca(OH) 2 

or by the reduction of a soluble arsenic compound, jirsenious 
acid, H 3 AsQ 3 for example, by zinc or magnesium, in the presence 
of hydrochloric or sulfuric acids, 

H 3 As0 3 + 3Zn + 6HC1 = AsH 3 + 3ZnCl 2 + 3H 2 O 

In this case, a large amount of hydrogen is given off, but the 
arsine is easily obtained in the pure state by passing the mixture 
through a tube surrounded by liquid air. The arsine condenses 
to a solid, melting at 119 and boiling at 55, while the 
hydrogen is not liquefied at this temperature. 

Arsine, either pure or mixed with hydrogen, may be almost 
completely decomposed into its elements by passing the gas 
through a red hot glass tube. The arsenic is desposited beyond 
the burner as a black mirror. Arsine kindles easily and burns 
with a pale blue flame, producing water and a white smoke of 
arsenic trioxide. Within the flame itself the arsenic is evidently 
in the free state, because if a piece of porcelain or other cold 
object be held in the flame, a mirror of arsenic will be deposited 
upon it. The mirror of arsenic formed on the cold porcelain, or 
deposited in the glass tube beyond the burner, is so very easily 


seen, even when only the merest trace of a soluble compound of 
arsenic is introduced into the hydrogen generator, that this test, 
which is known as Marsh's test, is an exceedingly valuable one. 
Antimony behaves in much the same way, but the two mirrors 
may be easily distinguished as will be seen. The zinc, acid and even 
the glass used must be especially pure and a " blank" experiment, 
run without the addition of the solution suspected to contain the 
arsenic, to establish their purity, must be carried out. The test 
is almost too delicate for ordinary work. 

Arsine is exceedingly poisonous, and the inhalation of even 
very small quantities has produced death. 

Arsenic Trioxide. Arsenic trioxide, or simply arsenic, as it is 
usually called, is the most important compound of arsenic. It 
occurs in nature in small quantities, but is generally obtained by 
the sublimation of the flue dust from smelters which roast arsenic 

There are two different crystalline forms of the trioxide, 
octahedral and monoclinic, and also an amorphous modification. 
In the process of purification by sublimation the arsenic trioxide 
collects as an amorphous glass-like substance, which, upon being 
kept for some time, gradually becomes milk white and looks 
very much like porcelain. The change in appearance is due to 
the passage of the amorphous into a crystalline (octahedral) 
modification. It is not very soluble in water, (about 2 grm. 
per 100 of water at 25). As usual the amorphous is more soluble 
than the crystalline. 

The vapor density of the trioxide at temperatures near 600 
corresponds to the formula As 4 6 and on this account it is some- 
times called the hexoxide, but at 1,800 the formula indicated is 
As 2 3 , and this simpler one is that generally used. 

The trioxide is very easily reduced by charcoal or by potassium 
cyanide with the formation of the metal. It acts both as an 
acid- and as a base-forming oxide, but is only slightly basic. 
With concentrated sulfuric or hydrochloric acid, it forms compli- 
cated sulfates or the trichloride, AsCls. These salts are com- 
pletely hydrolyzed upon dilution with much water. The tri- 
oxide is highly poisonous, but if taken in repeated small doses, 
the system becomes accustomed to the substance and quantities 
may be taken with safety which would ordinarily be fatal. 

GROUP V 431 

Arsenious Acid. A solution of arsenic trioxide in water is 
slightly acid, and is called arsenious acid. When attempts are 
made to obtain the acid in the anhydrous state, it decomposes 
into water and the trioxide. The trioxide is easily soluble in 
sodium hydroxide solution, and from this sodium arsenite, 
N^AaOojijw frp. oMfl.inpd. This sodium arsenite is an~ 

~~C^_ ^ JfcA^i ______ V-" mi I 

tant article of commerce and is used for making sheep "dips" as 
well as for manufacturing and chemical purposes. Like all other 
arsenic compounds, it is very poisonous. The arsenites of most 
of the metals, other than the alkalies are but slightly soluble. 
London purple, impure calcium arsenite and arsenate obtained 
as a byproduct in the manufacture of analine dyes, and Paris 
green, a copper acetate and arsenite, CugCC^Ha 

used as insecticides, and Scheele's green, CuMAsO 3 , is sometimes, 
though rarely at the present day, used as a pigment. It is 
objectionable on account of its poisonous nature. Ferric hydrox- 
ide forms very difficultly soluble arsenites, and is used as 'an 
antidote in cases of poisoning with arsenic trioxide or its salts. 
The arsenites are easily oxidized to the arsenates. 

The Halogen Compounds. Arsenic trichloride, AsCls, is the 
most important halogen compound of arsenic. It is a colorless, 
oily liquid, boiling at 130. It may be formed by the action of 
chlorine upon arsenic, or by hydrochloric acid upon the trioxide. 
When dissolved in water, it undergoes very extensive hydrolysis 
with the formation of hydrochloric acid and arsenious acid, which 
in turn is largely decomposed into the slightly soluble trioxide and 
water. This action is much like that between phosphorous tri- 
chloride and water, with the exception that the hydrolysis of the 
arsenic compound is reversed by the addition of concentrated 
hydrochloric acid. Because of this, the trioxide is more soluble 
in hydrochloric acid than in water. Because of the volatility of 
arsenic trichloride, precautions have to be taken before the 
evaporation of an acid solution of arsenic containing chlorine as 
ion to see that the arsenic is oxidized to the pentavalent form, 
and that no reduction takes place during the process. Arsenic 
trifluoride is a colorless liquid, while the bromide and iodide are 

Arsenic Trisulfide. Arsenic trisulfide, As 2 S 3 , is found in nature 
in the mineral called orpiment, which, upon being ground, yields 


a lustrous yellow powder, which was formerly used as a pigment. 
It may be prepared in the laboratory by passing hydrogen sulfide 
through an acid solution of an arsenite, or of arsenious acid. 
Under these conditions, it comes down promptly as a flocculent 
yellow precipitate. If the hydrogen sulfide is passed through a 
pure solution of arsenious acid, no precipitate will appear, but 
the solution becomes yellow in color, and contains the sulfide in 
colloidal solution in the form of very small particles, as is shown 
by the fact that the path of a ray of light through the solution is 
visible, and the solution polarizes light. The addition of acids 
and of neutral salts will cause the particles to gather into flocks 
and to precipitate. In bringing about this result, the salts are 
more active the higher the valence of the cation, but the anion is 
almost without influence. 

The sulfide is practically insoluble in even the strongest cold 
hydrochloric acid, but it is slowly dissolved by boiling concen- 
trated hydrochloric acid with a volatilization of arsenic trichlo- 
ride. Strong HC1 and KC1O 3 will dissolve it with the formation 
of arsenic acid, H 3 As04, sulfuric acid and reduction products of 
the chlorate. It is soluble in the alkali sulfides, giving thio- 
arsenites (NH 4 ) 3 AsS 3 , if the sulfide is colorless; or thioarsenates, 
(NH 4 ) 3 AsS 4 , if a polysulfide is used. The equations are 

As 2 S 3 + 3(NH 4 ) 2 S = 2(NH 4 ) 3 AsS 3 
2As 2 S 3 + (NH 4 ) 2 S 5 = 2As 2 S 5 + (NH 4 ) 2 S 
As 2 S 5 + 3(NH 4 ) 2 S = 2(NH 4 ) 3 AsS 4 

The corresponding sulfides of sodium and potassium will, of 
course, give sodium or potassium thioarsenites, or thioarsenates. 
The extra sulfur of the polysulfide is the oxidizing agent, which 
changes the trisulfide to the penta-, which is then dissolved to 
form the thioarsenate. When an acid is added to one of these 
solutions, thioarsenious or thioarsenic acid is formed. Each 
of these at once decomposes evolving hydrogen sulfide and 
precipitating a trisulfide from the thioarsenious acid, or the 
pentasulnde, As 2 S 5 , from the thioarsenic acid. This behavior 
of the sulfide toward alkali sulfides is very much like that of the 
sulfides of tin and puts arsenic and tin into the same analytical 

GROUP V 433 

Colloidal Solutions. Colloidal solutions have been repeatedly 
mentioned, and in each case it has been . pointed out that they 
are visibly heterogeneous when a powerful beam of light passing 
through them is viewed by a microscope, at right angles to the 
light. The minute particles in suspension are seen as points of 
light in a ceaseless dance called the Brownian movement. As 
was explained on p. 75, this movement is supposed to be due to 
the hammering of the molecules of the liquid upon the suspended 
solute. Every solid or non-miscible liquid when exposed to water 
takes on an electric charge either positive or negative depending 
upon which kind of ion it adsorbs in the greater amount. Most 
colloids become negatively charged, but the hydroxides of the 
metals are positive. These charges on the colloid help to hold 
the particles apart and to keep the solute in suspension. If in 
any way these charges are neutralized, the particles collect to- 
gether and soon settle out as a precipitate. 

Arsenic sulfide is a negatively charged colloid, and the neutral 
salts precipitate it because the sulfide adsorbs enough of the cation 
to neutralize its negative charge; and then the particles gather 
together like newly churned butter until they are large enough to 
filter out. 

It will be remembered that colloids are almost without effect 
upon the boiling- and freezing-point of the solvent. 


Arsenic Pentoxide and Arsenic Acid. When arsenic or the 
trioxide is treated with concentrated nitric acid, a solution is 
obtained from which, after high concentration, a hemihydrate of 
orthoarsenic acid, 2H 3 As04-H 2 O, crystallizes. This acid is a 
white very soluble and deliquescent substance. Like phosphoric 
acid, to which it is analogous, it is tribasic and forms three series 
of salts. The soluble normal salts are like the corresponding 
phosphates in that they are largely hydrolyzed and their solutions 
are highly alkaline. When the orthoarsenic acid is heated, it 
loses water and passes into the pentoxide, As 2 Os. When ortho- 
phosphoric acid is heated, the loss of water stops with the forma- 
tion of metaphosphoric acid and does not go so far as to form the 



The orthophosphates and the orthoarsenates are isomorphous 
and have about the same solubilities, so that it is not altogether 
an easy matter to distinguish between the two groups of salts. 
The sharpest distinction rests upon the fact that arsenic forms a 
difficultly soluble sulfide with hydrogen sulfide in acid solutions, 
while phosphorus does not. 

Sodium arsenate r Na 2 HAsQ4,j and the bi- or pyroarsenate, 
Na 4 As 2 07, have become important articles of commerce, being 
used to prepare the lead arsenates, Pb 3 (AsO 4 ) 2 or PbHAsO4 
which are extensively employed as sprays to protect fruit trees 
from the attacks of insects. 

Pentasulfide. Arsenic pentasulfide is a yellow powder, which 
may be obtained, as mentioned above, by acidulating a solution of 
a thioarsenate. It is also prepared by passing a rapid stream of 
hydrogen sulfide into a solution of arsenic acid in concentrated 
hydrochloric acid. It is a rather unstable substance, readily 
decomposing into the trisulfide and sulfur and having about the 
same solubilities as the trisulfide. In dissolving in the alkali 
sulfides, thioarsenates are formed whether the sulfide be colorless 
or yellow. 

Divalent Compounds. The best known of these compounds 
are the iodide, AsI 2 and the sulfide, As 2 S 2 . The latter occurs in 
nature and is known as realgar. It has a red color, melts easily, 
and burns to arsenic trioxide and sulfur dioxide. It may be 
made by heating together arsenical pyrite and pyrite, 

2FeAsS + 2FeS 2 = 4FeS + As 2 S 2 


Occurrence. Very small quantities of antimony are found 
free in nature. The principal ore is a trisulfide or stibnite, 
Sb 2 S 3 . From stibnite the metal is prepared either by fusion 
with metallic iron, which yields antimony and ferrous sulfide, or 
the ore is first roasted, forming sulfur dioxide and antimony 
tetroxide. SbsO*: the latter is then reduced by carbon. The 
metal prepared by either of these processes is not pure and must 
be remelted with sodium sulfate and slag from previous opera- 
tions and finally with potassium carbonate and slag. 

GROUP V 435 

Properties. When pure, antimony is a silvery white, highly 
crystalline metal, which may be easily powdered and which has 
a density of 6.7. It melts at 630 and boils at 1,440. The 
vapor density indicates a mixture of Sb 2 and Sb 4 in the vapor at 
1,634. Besides the ordinary form which has just been briefly 
described, there is a very unstable yellow modification, cor- 
responding to yellow phosphorus, a more stable black form, 
and an explosive kind which is deposited by electrolysis from 
concentrated solutions of the trichloride. Antimony is a con- 
stituent of several important alloys. Type metal contains two 
parts of lead and one part each of tin and antimony. Like most 
of the alloys of antimony, it expands upon solidification and 
hence makes sharp castings. Babbit metal contains 1.5 per cent, 
copper, 13 per cent, antimony, 45.5 per cent, tin and 40 per 
cent. lead. It is used as an antifriction metal for the bearings 
of machinery. 

Antimony is stable in the air at ordinary temperatures, but 
burns when heated, forming the trioxide or tetroxide, Sb 2 4 . 
It combines directly with the halogens and is oxidized by nitric 
acid to either the trioxide or antimonic acid, HsSbO^ with the 
formation of nitrogen peroxide. Hot concentrated sulfuric acid 
forms the sulfate Sb 2 (S04)3 and sulfur dioxide. Dilute sulfuric 
acid and hydrochloric acid are without action upon the pure 

Antimony forms two series of compounds; in the one it is 
trivalent, and in the other pentavalent. The trivalent are the 
more numerous and important. 


Stibine. One strong point of resemblance between antimony 
and the preceding members of the group is the formation of the 
hydrogen compound, stibine, SbH 3 . This is a colorless, gaseous 
substance, liquefying at 18 and freezing at 91.5. It is 
very poisonous. It is formed by the action of zinc in acid solu- 
tion upon a soluble antimony compound. The reaction is very 
similar to that for the preparation of arsine. Stibine is even 
more readily decomposed than arsine and may explode. When 
passed through a hot tube, the decomposition goes on rapidly 


at 200 and the antimony is deposited as a fusible non-volatile 
mirror on both sides of the heated portions of the tube. The 
mirror of arsenic from arsine is always deposited beyond the hot 
part of the tube and volatilizes easily without melting. 

The Trioxide and Its Acids. The density of the vapor of 
antimony trioxide at 1,560 corresponds with the formula, Sb 4 O 6 , 
analogous to that of phosphorus or arsenic trioxides at lower 
temperatures. In spite of this the formula of the substance is 
usually written Sb 2 O 3 , and it is called the trioxide, rather than 
the hexoxide. It may be obtained by acting upon antimony 
with nitric acid or by heating the metal in a limited supply of 
air. It is a white crystalline substance which is practically 
insoluble in water, dilute sulfuric or nitric acid; it is soluble in 
moderately dilute hydrochloric or tartaric acid and in concen- 
trated sulfuric acid and in solutions of sodium or potassium hy- 
droxide. These properties show that it is both basic and acidic, 
and therefore not strong in either way. Both series of salts are 
largely hydrolyzed. 

The hydroxide, Sb(OH) 3 is a white solid which has the same 
solubilities as the trioxide. When the oxide or hydroxide dis- 
solves in sodium hydroxide, sodium meta-antimonite, NaSbO 2 , 
is formed. Antimony hydroxide is soluble in potassium hydro- 
gen tartrate KHCJET^Ofi or " cream of tartar," as it is often 
called, owing to the formation of the complex salt tartar emetic, 
or potassium antimonyl tartrate, K(SbO)C4H 4 O 6 'J^H 2 O. This 
salt is important in medicine and also in the laboratory, since it 
is almost the only antimony compound which will form a clear, 
neutral solution. The explanation of this is found in the very 
small ionization of the ion, Sb +++ from the complex antimonyl 
tartrate ion, SbOCJ^Oe". The group SbO is called antimonyl 
and acts like a monovalent cation. The solubility of the trioxide 
in tartaric acid is due to the formation of antimonyl tartrate ion. 
Because of the weakness of the basic properties of the hydroxide, 
no carbonate is known and the chloride and sulfate are largely 
hydrolyzed with the precipitation of basic salts. When these 
basic salts are boiled with water, the hydrolysis becomes com- 
pleted and the trioxide is formed. 

Halogen Compounds. The most important halogen compound 
is the trichloride, SbCls. This may be readily prepared by dis- 

GROUP V 437 

solving the trisulfide in strong hydrochloric acid and distilling. 
The water passes off first and then the trichloride. It is a white 
solid, melting at 72.0 and boiling at 233. With a small quan- 
tity of water it gives a white precipitate of the oxy chloride, 
SbOCl, which is sometimes called antimonyl chloride. With 
more water the compound Sb 4 5 Cl 2 is formed, and when this is 
washed with hot water the oxide is left behind. Hydrochloric 
acid will reverse this hydrolysis even more readily than that 
of arsenic trichloride. 

The trichloride is used in medicine as a caustic and technically 
for bronzing iron, and in dyeing as a mordant. 

The trifluoride, SbF 3 forms colorless crystals, which are de- 
liquescent and whose solution may be diluted without precipita- 
tion, owing probably to slight ionization of the salt. 

The tribromide, SbBr 3 , and triiodide SbI 3 are formed in the 
same way as the chloride. They are easily decomposed by water. 

Other Salts. The salts of the oxyacids in general are not very 
stable and are hydrolyzed even more readily than the halogen 
salts. The sulfate, Sb 2 (S0 4 ) 3 , and the nitrate, Sb(N0 3 ) 3 , may be 
obtained in solution, but the latter will not stand dilution and 
soon decomposes on standing, in spite of an excess of acid. 

Antimony Trisulfide. Antimony trisulfide, Sb 2 S 3 is known 
both in the crystalline and amorphous state. The crystalline 
modification is black, melts at 450, and has a metallic luster. 
It is found in nature and is known as stibnite; it may also be 
formed by heating the red amorphous form to 220. The 
amorphous modification is precipitated when hydrogen sulfide 
is passed into an acid solution of an antimony salt. It is orange 
red in color and is soluble in concentrated hydrochloric acid, but 
is precipitated upon dilution. A method for separating arsenic 
and antimony may be founded upon the difference in the behavior 
of their sulfides toward hydrochloric acid. Arsenic sulfide is 
completely precipitated in very strong hydrochloric acid, while 
antimony is unaffected. The solution is then filtered and diluted, 
when, upon the passage of more hydrogen sulfide, antimony is 
quantitatively precipitated as the sulfide. 

Antimony trisulfide is soluble in solutions of sodium or am- 
monium sulfide, or poly sulfide, with the formation of thio- 
antimonites with the sulfides, and thioantimonates with the 


polysulfides, the reactions being strictly analogous to those for 
the arsenic compound; 

Sb 2 S 3 + 3(NH 4 ) 2 S = 2(NH 4 ) 3 SbS 3 
2Sb 2 S 3 + (NH 4 ) 2 SS 4 = 2Sb 2 S 6 + (NH 4 ) 2 S 
Sb 2 S 5 + 3(NH 4 ) 2 S = 2(NH 4 ) 3 SbS 4 

The most common thioantimonate is that known as Schlippe's 
salt, Na 3 SbS 4 -9H 2 O. 

When the solutions of the tiro salts are acidulated, thio- 
antimonious, or thioantimonic acids are formed. These acids 
are unstable and promptly break down into Hydrogen sulfide 
which escapes as a gas, and antimony tri- or pentasulfides, which 
are reprecipitated. The solubility of the sulfide in ammonium 
sulfide and its reprecipitation by acids places antimony in the 
analytical group with tin and arsenic. 


The pentavalent compounds are formed from the trivalent by 
the action of oxidizing agents. 

Antimony Pentasulfide. The method for the formation of the 
pentasulfide has just been given in connection with the trisulfide. 
The oxidizing agent in this case is the polysulfide. The penta- 
sulfide is a yellow-red substance which like the trisulfide is 
soluble in strong hydrochloric acid, the sulfides of ammonium, 
and of the alkalies. It is reprecipitated from the hydrochloric 
acid solution on dilution, and from the other solutions upon 
acidulation. The thioantimonates have already been mentioned 
in connection with the trisulfide. 

Antimony Pentachloride. Antimony pentachloride, SbCl 5 , is 
a colorless liquid freezing at 6. It is formed when an excess 
of chlorine acts upon powdered antimpny or by the action of 
chlorine upon the trichloride. 

When the pentachloride is treated with a little water, various 
hydrates are formed, while with more water it is completely 
hydrolyzed, giving antimonic acid, or its anhydride, the pentoxide. 

Antimonic Acids and the Pentoxide. As was the case with 
phosphoric and arsenic acid, there are three antimonic acids, 
Qrthoantimonic, H 3 SbO 4 , pyroantimonic^H 4 Sb 2 O7, and metanti- 

GROUP V 439 

monic, HSbOs acids. When these acids are heated they pass 
into the pentoxide and water. The acids and also the pentoxide 
are sparingly soluble in water or acids, but dissolve in potas- 
sium hydroxide. These properties indicate that the pentoxide is 
a purely acid-forming oxide. Potassium metantimonate, KSbO 3 , 
is formed by fusing potassium nitrate with powdered antimony. 
When boiled with water it forms the acid pyro salt, K2H2Sb2O7, 

2KSb0 3 + H 2 O = K 2 H 2 Sb 2 O7 

This potassium salt is easily soluble in water, but the correspond- 
ing sodium salt requires 350 parts of water for its solution, and 
hence is precipitated upon the addition of a concentrated solution 
of the potassium salt to a strong solution of a sodium salt. This 
is interesting as it is one of the least soluble inorganic sodium 
compounds. As a test for sodium, it is far less sensitive and 
characteristic than the flame-test. 

Antimony tetroxide, Sb 2 O 4 , is formed by heating the trioxide 
to somewhere between 400 and 775 in the air. In the absence 
of reducing agents it is stable up to high temperatures. It is 
neither an acid nor a base. One large use for the substance is in 
the manufacture of enamel ironware for cooking purposes. In 
view of the highly poisonous character of antimony compounds, 
this use would Hot seem to be' wholly free from objection. 


General. Bismuth is far more metallic in its character than 
the other members of the group. The trivalent compounds 
are the only stable and important ones. It does not form a 
compound with hydrogen. Bismuth trihydroxide, Bi(OH) 3 is 
the strongest base in the family. It does not act as an acid, 
and does form carbonates, nitrates, phosphates, sulfates, etc. 
These salts are hydrolyzed by water so the hydroxide is not a 
really strong base. The trisulfide does not form thio salts with 
the alkali sulfides. 

Occurrence and Preparation. Bismuth is found free in nature 
and to a smaller extent as the sulfide, Bi 2 S 3 , and oxide, Bi 2 3 . It 
is obtained from its ores by first roasting them to remove sulfur, 
and then heating in a crucible or reverberatory furnace, with 


carbon to act as a reducing agent, iron to reduce unchanged 
sulfide, and to combine with arsenic, and fluxing material to form 
an easily fusible slag. 

Properties. Bismuth is a hard, brittle, crystalline metal with 
a high luster and a distinct reddish tinge of color. It has a 
density of 9.75, melts at 268 and boils at 1,420. Its molecular 
weight lies between that required for the formulas Bi and Bi 2 . 
The atomic weight is 208. Several alloys of bismuth, lead, tin 
and cadmium have been prepared, which have a melting-point 
below the boiling-point of water. Lipowitz metal containing 15 
parts bismuth, 8 parts lead, 4 parts tin, and 3 parts cadmium, 
melts at 60. Higher melting alloys are used as safety plugs in 
steam boilers and in making automatic fire extinguishers. Bis- 
muth alloys expand on solidification and make good castings. 
Some of them are used as stereotype metal. 

The metal is but very slightly acted upon in the air at ordinary 
temperatures, but when heated it is slowly oxidized to the tri- 
oxide. It requires a fairly strong oxidizing agent to transform 
it into ion, so it is not dissolved by hydrochloric acid in the 
absence of air. It is dissolved by hot concentrated sulfuric acid 
with the evolution of sulfur dioxide and also in the cold by nitric 
acid or aqua regia. In every case, a salt of trivalent bismuth is 

Bismuth Trioxide. Bismuth trioxide, Bi 2 O 3 , is a yellow sub- 
stance and is found in nature as bismuth ocher. It is made in 
the laboratory by heating the hydroxide, carbonate or nitrate. 
The trioxide has the same solubility in acids as the hydroxide, 
which is given below, and forms the same salts. It is not soluble 
in bases. The hydroxide, Bi(OH) 3 is a white amorphous powder 
and is prepared by precipitating a bismuth salt with ammonium 
or sodium hydroxide. It is not soluble in excess of the base and 
in this way differs radically from antimony hydroxide. Acids 
dissolve the hydroxide with the formation of salts, such as the 
trichloride, BiCl 3 , or nitrate, Bi(NO 3 ) 3 . When solutions of the 
salts are diluted; basic salts are precipitated, such as the oxy- 
chloride, BiOCl. In fact, clear solutions of bismuth salts can be 
prepared only in the presence of an excess of acid. 

Bismuth Salts. The nitrate, Bi(N0 3 ) 3 -5H 2 0, is the most 
important salt. It forms clear colorless crystals which are 

GROUP V 441 

soluble in dilute nitric acid, but upon dilution the solution de- 
posits the basic salt, BiON0 3 . This is called bismuth subnitrate 
and is largely used in medicine. The trichloride, BiCl 3 -H 2 0, is 
colorless and is soluble in dilute hydrochloric acid. It forms the 
very slightly soluble oxychloride upon dilution, which is so diffi- 
cultly soluble that it is used in the identification and separation 
of bismuth. The bromide and iodide are much like the chloride. 
In general, all the bismuth salts are easily hydrolyzed forming 
basic salts. This hydrolysis may be reversed by acids. 

Bismuth Trisulfide. Bismuth trisulfide, Bi 2 S 3 , is formed by 
the direct union of the elements, or by the action of hydrogen 
sulfide upon a moderately acid solution of a bismuth salt. It is 
brownish-black and is not soluble in water, cold dilute acids, the 
soluble sulfides, or polysulfid.es. It is soluble in concentrated 
hydrochloric and hot dilute nitric acid; in the latter case the 
sulfur is either liberated or oxidized to the sulfate. These prop- 
erties place bismuth in the analytical group containing mercuric 
mercury, lead, copper, and cadmium. It is distinguished from 
mercury by the fact that its sulfide is soluble in hot dilute nitric 
acid, while that of mercury is not; from lead by the slight solu- 
bility of lead sulfate in dilute sulfuric acid; and from copper and 
cadmium by the fact that the hydroxide of bismuth is not soluble 
in an excess of ammonium hydroxide, while those of the other 
two metals are. It is finally identified by dissolving a little of 
the hydroxide in a few drops of dilute hydrochloric acid and pour- 
ing the solution into water, when a white precipitate of the 
oxychloride will form. 

Other Compounds. Besides the trioxide, bismuth forms the 
monoxide, BiO, the tetroxide, Bi 2 O 4 , and the pentoxide, Bi 2 5 . 
A monosulfide, BiS, analogous to the monoxide, is known. Bis- 
muth dichloride and dibromide have been prepared, but no penta- 
chloride or bromide. 

The facts collected in the following table, show that nitro- 
gen and bismuth differ more from the other three elements than 
these do from one another, but that the elements taken as a 
whole form a well-defined family with fairly regular changes in 
properties as the atomic weight is increased. Some of the 
more striking of these changes may be summarized as follows: 
As the atomic weight increases, the elements change from non- 



metals to metals; the melting-point, boiling-point, and density 
of the element rises; the boiling-point of the hydrides (except 
ammonia) rises; the stability of the hydrides decreases; the 
basic properties of the hydrides decrease ; the acid properties 
of the oxides decrease; the basic properties of the oxides increase. 







Atomic wt 









AS4 AS2 

Sb4-Sb 2 


Density . 

- 210 
- 194 

+ 44.0 
1 8-2.3 

+ 480.0 
3 . 9-5 . 7 

1,440 app. 

1,420 app. 

Occurrence. . . . 


Cb. never 










B.p. of hydro com. 
Comb with acids 

- 33.5 


- 86 


- 55 

- 18 


A = any ; . . . . 


PI 2 

PA S * 



SbA B t 

BiCl 2 
BiF 6 

(ous) . 

NzO 1 





N2O8 1 


As4Oe 1>2 

Sb4O 6 1>2 

Bi 2 Oa 2 



P 2 O4 

Sb 2 O4 

Bi 2 O4 


N 2 O s i 

P2O* 1 

AS2OB 1 

Sb 2 O 6 1 

Bi 2 O B 






HAsO2 J 







HAsOs 3 



H 4 P 2 O7 
H 3 PO 

H4AS 2 O? 3 


H4Sb 2 O 7 


N 4 S 4 

P 2 Ss 

As 2 S2 

Sb 2 Sa 

Bi 2 S 2 
Bi 2 Sa 


N 2 S. 

P 2 S S 

As 2 S s 

Sb 2 S 8 

Bi 2 S 6 

* Except PI 6 which is unknown, 
t Except SbBrs which is unknown. 1 
1, 2 = both, acid and base forming. 3 

= acid forming. 2 = base forming. 
= known only in salts. 


U <- W <- Mo <- Cr <- O - S -> Se -* Te 

Chromium sub-group Oxygen sub-group 

The members of Group VI fall naturally into two well-marked 
families or sub-groups; the first is called the oxygen or sulfur 
group and is made up of oxygen, sulfur, selenium, and tellurium. 
These are non-metals, although tellurium posesses some of the 
characteristics of the metals. Their melting- and boiling-points 
are comparatively low and rise regularly with the atomic 
weights. These elements all form compounds with the metals 
which are closely analogous, that is to say the oxides, sulfides, 
selenides, and tellurides are much' alike. They all form com- 
pounds with hydrogen which are readily volatilized. ^ 

The members of the chromium family are chromium, molyb- 
denum, tungsten, and uranium. These are all distinctly metallic, 
have very high boiling-points, are exceedingly hard, and do not 
form hydrogen compounds analogous to water, hydrogen sulfide, 
etc. The two sub-groups then differ markedly in some respects 
and the points of dissimilarity are about what would be expected 
from what has been noticed in the other groups. The elements 
of the chromium sub-group are more metallic than those of the 
sulfur family, have greater densities, and higher melting-points. 
The members of the whole sixth group have this in common that, 
with the exception of oxygen, they each have a maximum normal 
valence of six and form the oxide ROs, in which R stands for any 
member of the group. Sulfur forms persulfuric acid, H 2 S2Os, in 
which its valence appears to be seven and possibly chromium has 
this same valence in the very unstable perchromic acid. These 
heptavalent compounds are to be regarded as rather abnormal 
for the group. The oxide is in each case the anhydride of an acid. 
But the strength of the acid decreases as the atomic weight 
increases until with tellurium and uranium distinctly basic prop- 



erties appear. Sulfur, selenium, and tellurium form dioxides, 
S0 2 , Se0 2 , Te0 2 . The first two are acid-forming exclusively, 
the last is both acidic and basic. - Most of the members of the 
chromium group form the oxides RO, R 2 0s, R0 2 as well as 
RO 3 . These are in general base-forming, but are often both 
basic and acidic. 

The members of the oxygen-sulfur family have been discussed 
among the non-metals and it now remains to take up a part of 
the chemistry of the chromium sub-group. The members of this 
family form so many complex compounds that anything like a' 
detailed treatment in a book of this kind is out of the question. 


i (. 
General. Chromium is di-, tri-, hexa- and possibly hepta- 

valent, and has a very tangled relationship, because every time it 
changes its valence, it takes on points of similarity to different 
elements. In the hexavalent state, it is acid-forming, and 
resembles sulfur in sulfates; in the trivalent condition, it is 
very much like aluminum and ferric iron; when divalent it is 
much like ferrous iron. Chromium was given its name because 
all its compounds are highly colored. 

Occurrence. Chromium does not occur free in nature and its 
only important ore is chrome-ironstone which is a mixture of 
chromium trioxide, Cr 2 Os, and ferrous oxide, FeO, to which the 
formula Fe(CrO 2 ) 2 is usually given. This mineral is somewhat 
uncommon although by no means rare. 

Preparation. Pure chromium may be most readily prepared 
by the Goldschmidt process, reducing the trioxide with powdered 
aluminum. The oxide is readily reduced by carbon at the 
temperature of the electric funace, but the products contain 
some carbon. An alloy of chromium and iron containing 60 
per cent, of the former known as ferro chrome, is made on a 
large scale in electric furnaces by the reduction of chrome-iron- 
stone. Ferro chrome is used in the preparation of chromium 
alloy steels which are very hard. 

Properties. Chromium is a tin white, very hard and brittle 
metal melting at about 1,520. It may be distilled at the tem- 
perature of the electric arc furnace. The metal is apparently 


not attacked by the air at ordinary temperature, but burns bril- 
liantly at the temperature of the oxyhydrogen blow-pipe. It dis- 
solves in cold concentrated or hot dilute hydrochloric, sulfuric, 
and oxalic acids, but is not attacked by nitric, chromic, chloric, 
or perchloric aicds. This is surprising, because these latter acids 
are more active oxidizing agents than the others, and curiously 
enough a piece of chromium that has been in nitric acid will not 
dissolve when placed in hydrochloric acid. This is expressed by 
saying that chromium becomes passive in the oxidizing acids, and 
that this passivity persists when it is placed in other acids. Con- 
tact with the air or with free chlorine or bromine as well as with 
the oxidizing acids will make chromium passive, while reducing 
agents in general will render it active. For example, if a piece 
of passive chromium in hydrochloric acid be touched with a stick 
of zinc, cadmium, copper or even active chromium, it will at once 
begin to dissolve with the evolution of hydrogen. There is a 
difference of potential of about 1.6 volts between the active and 
the passive forms. 

This change back and forth from active to passive in the 
presence of reducing or oxidizing agents suggests that the passiv- 
ity may be due to a layer of oxide; but this explanation is not 
entirely free from objection. This passivity of chromium is per- 
haps connected with the fact that certain of the chromium alloys 
with iron and other metals are so resistant to chemical action 
that they may with advantage be used in place of platinum for 
many chemical purposes. A chromium anode dissolves either 
as the chromous or chromate ion according to the current 

Chromous Compounds. The divalent or chromous compounds 
may be prepared by dissolving the metal in acids; by the reduc- 
tion of the chromic salts with metallic zinc ; or by heating chromic 
chloride in a stream of hydrogen. In the solid state they are red,, 
blue, or yellow, but in solution are always blue which seems 
to be the color of the chromous ion, Cr ++ . They change so 
easily to the trivalent or chromic compounds, that they will 
even be oxidized by water with the evolution of gaseous hydrogen, 
and are therefore very unstable. The principal ones are the 
hydroxide, Cr(OH) 2 , chloride, CrCl 2 , sulfate, CrSO 4 -7H 2 0, 
and acetate, Cr(C 2 H 3 O 2 )2. The latter is a dark red crystalline 


salt which is not very soluble in water, but is dissolved by strong 
acids. It is about the most stable of the chromous compounds. 

Chromic Compounds. Chromium in the chromic compounds 
is trivalent, and the properties of the chromic ion are very much 
like those of aluminum, and the trivalent iron and manganese 
ions. All the chromic compounds are highly colored ; in solution 
most of them are violet and this is said to be the color of the 
chromic ion Cr + + + . Some solutions are green, but these have 
different chemical properties from the violet, and apparently 
contain complex ions. Chromium hydroxide, Cr(OH) 3 , is very 
difficultly soluble, and is precipitated when a soluble base is 
added to a chromic solution. 

Like aluminum hydroxide, it acts as a base and as an acid, and 
so dissociates as follows: 

Cr+ + + + 30H- <= Cr(OH) 3 * H+ + Cr0 2 ~ + H 2 O 

It is soluble both in acids and in excess of sodium or potassium 
hydroxide, forming in the first case chromic salts, and in the 
second chromites such as NaCrO 2 . Upon standing or more 
quickly when boiled, a green partially dehydrated form of the 
hydroxide is deposited from the chromite solutions. 

Ammonium hydroxide is too weak a base to dissolve chromium 
hydroxide with the formation of chromite, but it does dissolve it 
slightly owing to the formation of complex ammonia ions. 
Chromic hydroxide is a very weak base, and its salts, especially 
those of the weaker acids, are largely hydrolyzed, so the soluble 
sulfides precipitate chromic hydroxide instead of the sulfide. 
Carbonates usually precipitate the hydroxide, but under some 
conditions basic carbonates may be obtained. This action is 
much like that of aluminum salts. 

Chromic oxide^Cr 2 3j /may be made by heating the hydroxide 
or ammonium dichromate (NH 4 ) 2 Cr 2 07. 

(NH 4 ) 2 Cr 2 O 7 = Cr 2 3 + 4H 2 O + N 2 

and also by igniting a mixture of potassium dichromate, K 2 Cr 2 O7, 
and sulfur 

K 2 Cr 2 O 7 + S <=i Cr 2 3 + K 2 S0 4 
The oxide has a fine green color and is used as a pigment* 


Chromic Chloride. Anhydrous chromic chloride is made by 
heating the metal in a stream of chlorine. It sublimes as peach- 
blossom-colored scales which are so slowly soluble in water that 
they are apparently insoluble; but in the presence of minute 
quantities of a powerful reducing agent, for example, a very little 
chromous chloride, solution begins and goes on so rapidly that 
the temperature of the system is raised several degrees, and a 
large amount of the chloride goes in solution, where it exists in the 
chromic state. Alternate reduction to the chromous and oxida- 
tion to the chromic suggests itself as an explanation for the 
action of the reducing agent. 

The solution obtained in this way is green, and upon the addi- 
tion of silver nitrate, from one- to two-thirds only of the chlorine 
is precipitated as silver chloride. The remainder is evidently 
present in complex ions. When a dilute solution of the chloride 
is allowed to stand for some time, it becomes violet; and then 
the whole of the chlorine may be precipitated by silver nitrate. 
The violet solution is also a better conductor of electricity than 
the green. By carefully evaporating a solution of the chloride, 
green crystals of a hexahydrate, CrCl 3 -6H 2 O, may be obtained; 
under other conditions, two more modifications, one blue and the 
other green, each containing six molecules of water, may be pre- 
pared. When a hexahydrate is heated, it decomposes into the 
oxide Cr 2 Os, hydrogen chloride and water. 

Chromic Sulfate. The normal sulfate, Cr 2 (S04) 3 '15H 2 0, is 
violet in color, and its solution contains both the chromic and 
sulfate ions; but when heated, either dry or in solution, it turns 
green, and under certain conditions, passes into a state in which 
it will give reactions for neither chromic nor sulfate ions. Upon 
allowing the solution to stand at ordinary temperature, it slowly 
passes back into the violet modification. The electrical con- 
ductivity of the violet solution is much greater than that of the 
green under like conditions. 

Chromic sulfate combines with potassium or ammonium 
sulfate to form the corresponding alum, KCr(SO4)2'12H 2 O 
or NH 4 Cr(SO4) 2 -12H 2 O. 

These are violet in color and are isomorphs with ordinary 
alum. When a solution of potassium chrome alum is boiled, 
it turns green and will then refuse to deposit crystals upon 


cooling until it has stood for some days and changed back into 
the violet. 

Chromic Acid and the Chromates. Strong oxidizing agents in 
the presence of the alkalies or their carbonates will change the 
di- or trivalent chromium to hexavalent with the formation of 
chromates. Of these, perhaps, the potassium salt, K 2 CrC>4, is 
the most important although the sodium salt is cheaper and very 
largely used. Potassium chromate is made by heating a mixture 
of chrome-ironstone, potassium carbonate and limestone in 
contact with the air. The equation is 

4Fe(Cr0 2 ) 2 + 8K 2 C0 3 + 7O 2 = 8K 2 Cr0 4 + 8CO 2 + 2Fe 2 O 3 

The limestone is added to render the mixture porous and so 
assist the oxidation. The product is extracted with water, and 
any calcium chromate which may have formed is decomposed 
by the addition of potassium carbonate or sulphate. 

The sodium salt is prepared in the same way using the sodium 
instead of the potassium carbonate. The chromates are nearly 
ah 1 yellow in color and have about the same solubilities as the 
corresponding sulfates with which they are usually isomorphous. 
Lead chromate is very difficultly soluble, has a brilliant yellow 
color and is the pigment named chrome yellowy When attempts 
are made to prepare chromic acid, red crystals of chromic anhy- 
dride, CrOs, are obtained and apparently the acid, H 2 Cr04, 
cannot exist in the pure state. 

Bichromates. Chromic anhydride is made by adding con- 
centrated sulfuric acid to a solution of potassium dichromate; 
it crystallizes out and is washed with nitric acid. It will be 
recalled that sulfuric acid combined with sulfuric anhydride S0 3 
to form pyrosulfuric acid, H 2 S 2 C>7, and that this acid yields a 
series of salts called the pyrosulfates. In much the same way, 
chromic anhydride will combine with chromates forming the 
dichromates as they are, or pyrochromates, as they might be 

K 2 CrO 4 + Cr0 3 = K 2 Cr 2 O 7 

So easily does this reaction take place and so readily is chromic 
anhydride formed in solution, that dichromates are made by 
simply acidulating solutions of the chromates. This equation 
may best be written ionically: 


4K++ 2CrO 4 +2H+ + S0 4 --^4K+ + Cr 2 07~- + S0 4 

H 2 

or omitting the ions which are unchanged 

2CrO 4 --- + 2H+^Cr 2 7 - ' + H 2 O 

The reaction being reversible, it will at once be seen that acids 
by increasing the concentration of the H+ will cause it to go 
toward the right, and transform chromates into dichromates, 
while bases which will decrease the concentration of H + will 
cause dichromates to change into chromates. The above equa- 
tion will also make it clear why the difficultly soluble chromates 
such as barium or lead are precipitated from their soluble salts 
upon the addition of potassium dichromate. As fast as the chro- 
mate ion which always exists to a certain extent in any dichro- 
mate solution is used up, more is formed from the dichromate 
until the latter is finally practically all used up. The dichromate 
ion is orange red in color while the chromate is yellow; and the 
change of color may be used as a rough indicator of the presence 
of H+ or OH-. 

The tendency of chromic anhydride to combine with chromates 
is not exhausted with the formation of dichromate as is shown 
by the existence of such compounds as ammonium trichromate 
(NH 4 ) 2 Cr 3 Oio, and the tetrachromate (NH 4 ) 2 Cr 4 Oi 3 . 

This tendency for the formation of complex compounds is 
much more highly developed in molybdenum and tungsten. 

Chromates and dichromates are good oxidizing agents changing 
to the chromic compounds and giving up three positive charges in 
the transfer for each atomic weight of chromium. 

Dichromates in acid solution will oxidize alcohol to aldehyde; 
oxalate to carbon dioxide; ferrous salts to ferric; and liberate 
chlorine, bromine or iodine. In either acid or alkaline solution, 
they will change the sulfur ion to free sulfur. 

Chromyl Chloride. When a mixture of a dichromate and a 
chloride is treated with a concentrated sulfuric acid, a dark red 
fuming liquid distils over which bears a striking resemblance to 
bromine. This is chromyl chloride, CrOl^V 

4NaCl + Na 2 Cr 2 7 + 6H 2 SO 4 = 2CrO 2 Cl 2 +6NaHSO 4 +3H 2 O 



This chromyl chloride is analogous to sulfuryl chloride and like 
it is decomposed by water; 

2CrO 2 Cl 2 + 3H 2 O = H 2 Cr 2 O 7 + 4HC1 

From this, it is seen that chromyl chloride is an acid chloride 
but it is interesting to note that the corresponding uranium com- 
pound, UO2C1 2 , is a salt. This is in accord with the increase 
in the basic properties with the atomic weight. 

Photochemical Reactions. When gelatine or glue is treated 
with a compound of trivalent chromium, it is rendered insoluble 
in hot water. Use is made of this in tanning leather which con- 
tains a sort of gelatine. The chromates or dichromates do not 
act in this way, but if a mixture of a dichromate and glue is 
exposed to the sunlight, the dichromate is slowly reduced and 
then tans the glue. Several technically important photochemical 
processes are founded upon this. If a metal plate coated with a 
mixture of glue and ammonium dichromate is exposed under an 
especially prepared negative and then washed with hot water, 
the glue is washed away where the plate has been protected from 
the light, but not in the other portions. The metal plate is then 
treated with acids, and the parts not protected by the tanned 
glue are partially dissolved away. Such a plate may be inked 
and printed from in a regular printing press and the dark spots 
in the object will come dark in the print. 

Perchromic Acid. When hydrogen peroxide is added to an 
acid solution of dichromate, a very deep blue solution is ob- 
tained which soon evolves oxygen and turns green. The deep 
blue solution is supposed to contain per chromic acid which is 
thought to have the formula HCrO 4 . The appearance of this 
blue color is a delicate test for either hydrogen peroxide or a 

Analytical Reactions of Chromium. All the compounds of 
chromium are strongly colored; chromous, generally blue; 
chromic, violet or green; chromate, yellow, and dichromate 
orange red. Chromous compounds are too unstable to require 
much attention. Chromic compounds are so much like those 
of aluminum that they fall into the same analytical group, and 
are precipitated by ammonium hydroxide in the presence of 
ammonium salts. Chromium is most readily distinguished from 


iron and aluminum by the fact that its hydroxide is oxidized to 
sodium chromate in aqueous solution by sodium peroxide. The 
chromate may be recognized by its strong yellow color, by the 
formation of barium or lead chromate or by the blue color 
obtained on the addition of hydrogen peroxide to its acid solution. 


The rather rare element molybdenum is the second member of 
the chromium family. It does not occur free in nature and is 
found chiefly as molybdenite, the sulfide, MoS 2 , and also as 
wulfenite or lead molybdate, PbMoO4. When the sulfide is 
roasted in the air molybdenum trioxide, Mo0 3 , is formed. This 
may then be reduced to the metal by carbon in an electric 
furnace, by hydrogen at a high temperature, or by aluminum 
using the Goldschmidt process. It is a silvery white metal 
which is about as hard as iron. It resembles iron in that it can 
be forged when hot, takes up carbon and then can be tempered 
like steel. It is used in the form of ferro-molybdenum alloys 
in the manufacture of special alloy steels. The metal is scarcely 
affected by oxygen up to 600 when it forms the trioxide. It is 
insoluble in dilute acids, except nitric. 

Molybdenum forms three oxides: the sesquioxide, Mo 2 O 3 ; 
the dioxide, MoO 2 ; and the trioxide, MoO 3 . The first two 
are basic. The chlorides of molybdenum are Mo 3 Cl 6 , MoCl 3 , 
MoCl 4 , MoCl 5 . The trioxide is the anhydride of molybdic acid, 
H 2 MoC>4-H 2 O. This forms a series of molybdates which are simi- 
lar to the sulfates and the chromates not only in formula but also 
in solubility. Perhaps the most important of these salts, be- 
sides the naturally occurring lead compound, are the sodium 
and ammonium molybdates, Na 2 Mo04-10H 2 O and (NH 4 ) 2 MoO4. 
The latter is much used as a reagent. Molybdic acid and its 
salts have a great tendency to combine with the trioxide forming 
complex compounds and a very large number are known. The 
formula of one of the sodium salts will serve to show something 
of the complexity of these compounds, Na 2 MoioO 3 i-21H 2 O. 

Molybdenum trioxide will combine with other acids than 
molybdic acid to form complex acids. For example, it forms a 
series of such compounds with phosphoric acid, and these in turn 


form yellow crystalline very slightly soluble salts with ammo- 
nium, potassium, rubidium and cesium in the ionic state. Be- 
cause of this, when a dilute nitric acid solution of ammonium 
molybdate is added in excess to the solution of phosphate and 
the whole kept warm for some time to favor the formation of the 
complex phosphomolybdate ion, a yellow precipitate of ammo- 
nium phosphomolybdate, (NH^sPO^CMoC^n-GH^O, separates. 

The precipitate is soluble in excess of phosphoric acid or in 
alkalies, but is not soluble in dilute acids which implies that the 
phosphomolybdic acid is a strong acid. Arsenic acid and the 
arsenates react with the molybdates to form compounds which 
are very much like the phosphomolybdates in appearance and 

When an acid solution of a molybdate is acted upon by a 
reducing agent such as hydrogen sulfide, sulfur dioxide or dilute 
stannous chloride, a blue solution which deposits a blue pre- 
cipitate is formed. The composition of the precipitate is in 
doubt, but by some it is regarded as the molybdate of tetra- 
valent molybdenum, Mo(MoC>4)2 or MosOg. 


Tungsten is a rather rare element which has recently become of 
great economic importance. It occurs principally in wolfram, 
an iron-manganese tungstate (FeMn)WC>4. In the preparation 
of the metal from wolfram, the ore is first heated with sodium 
carbonate to form sodium tungstate which is then dissolved in 
water. When the solution of sodium tungstate is acidulated, 
tungstic acid is precipitated. By gently heating the acid, it is 
converted into the trioxide which may be reduced to the metal 
in a number of ways; by carbon, hydrogen, zinc, or aluminum at 
rather elevated temperatures. Because of the very high melting- 
point of tungsten which is about 3,500, the element is usually 
obtained in the form of a powder; but a homogeneous mass of 
the metal may be secured by reducing the oxide with aluminum, 
if an excess of powdered aluminum is used and the mixture is 
wet down with liquid air rich in oxygen before it is set off. The 
extra high temperature necessary for the fusion of the tungsten 
is furnished by the powdered aluminum burning in the oxygen 


of the liquid air. The almost instantaneous change of tempera- 
ture in this case from 182.5 to over 3,500 is, to say the least, 
striking. The metal is darker colored than zinc, is crystalline, 
harder than glass, very brittle, and has a density of 19.6. One 
of the modern applications of tungsten is in the construction of 
high efficiency incandescent electric lamps, a use for which its 
very high melting-point renders it well adapted. The prepara- 
tion of very fine wires required for such filaments is a matter of 
some difficulty, but it has been successfully done by compressing 
the powdered metal into rods and then rolling and drawing these 
down while they are heated by an electric current in an atmos- 
phere of hydrogen. Lamps prepared from this tungsten are 
very durable and because of their high efficiency have almost 
replaced the old carbon lamps. The old style standard 16-cp. 
carbon lamp consumed about 60 watts. The corresponding 
tungsten lamp uses 25 watts and gives about 20 cp., so the 
tungsten lamp gives one-fourth more light on less than half the 
energy which means that it converts the electrical energy more 
efficiently into light and produces less heat. It is estimated that 
the replacement of the carbon lamps by tungstens has made a 
saving of several hundred millions of dollars per year in the 
country's light bill. A still further improvement has been made 
by filling the lamps with nitrogen or still better, argon. In these 
0.5 of a watt or even a little less will give 1 candlepower. 

Another important application of tungsten is in the prepara- 
tion of special steels such as "high-speed steel." For this purpose 
a ferrotungsten is often prepared in the electric furnace. High- 
speed steel through increasing the rate at which machines 
and machinists can work, has lowered the cost of automobiles on 
an average of $200 per machine and made a corresponding 
saving in other lines. Hence tungsten is economically very 

The metal is fairly resistant to chemical action, but when 
powdered it will burn at a red heat to the trioxide. The ordinary 
acids act upon it very slowly; aqua regia dissolves it readily. 
It reacts with chlorine to form the hexachloride, WCle, from 
which the chlorides, WC1 2 , WC1 4 and WC1 5 , may be obtained 
by reduction. In addition to the trioxide already mentioned, a 
dioxide, WO2, is known. The sodium tungstate, Na 2 WO4'2H 2 O, 


is used as a mordant and for rendering cotton fabric fireproof, 
for which purpose it is dissolved in the starch with which the 
articles are treated. Very complex tungstic acids and the tung- 
states are formed by the union of the trioxide in varying quanti- 
ties with the acid or the tungstates. Compounds very similar 
to the phosphomolybdates are formed by the interaction of tung- 
stic acid with phosphoric and other acids which makes the chem- 
istry of tungsten very complex. 

The atomic weight of tungsten is 184.0. 


Uranium, which is the last element of the chromium family, 
has the highest atomic weight, 238.2, of any known element. 
It occurs in a number of rather rare minerals, but is chiefly ob- 
tained from pitchblende which contains the oxide, U 3 O 8 , or car- 
notite which is a vanadate of uranium and potassium. These 
ores are interesting as furnishing the principal source of radium 
compounds. The metal may be prepared from the tetrachloride 
by reduction with metallic sodium, or from the oxide by reduction 
with carbon or aluminum. Pure uranium is much like iron. 
It has a white color, takes a high polish, has a density of 18.7, 
and melts at about 1,500. The powdered metal is rather 
easily attacked chemically. It is oxidized in the air and burns 
in oxygen and the halogens at moderately elevated tempera- 
tures. There are two well-known oxides of uranium, the 
dioxide, U02, and trioxide, UO 3 . When treated with acids, 
the dioxide gives a series of salts called uranous salts such as 
uranous chloride, UC1 4 , uranous sulfate, U(SO 4 )2'8H 2 0, etc. 
The trioxide is both basic and acidic. With acids it forms a 
series of salts called uranyl salts. These contain the group 
UO2 + + which acts like a complex divalent cation. These salts 
have a yellow color and many of them have a greenish-yellow 
fluorescence. The most important of the uranyl salts and in 
fact the most important of the uranium compounds is uranyl 
nitrate, UO2(NO 3 )2-6H 2 O. The addition of a soluble hydroxide 
to a uranyl salt produces a precipitate of a diuranate, analogous 
to the dichromates. The sodium salt, Na2U207, is difficultly 
soluble and has a fine yellow color. It is an article of com- 


merce under the name uranium yellow and is used to color 
uranium glass which, like many other uranium preparations, has a 
greenish-yellow fluorescence. This phenomenon is due to the 
absorption of the invisible ultra-violet rays of the sunlight by 
the uranium compound and their conversion into visible greenish 
rays. Any substance which is able to become luminous, while 
under the excitation of some form of energy, without becoming 
heated to incandescence is said to be fluorescent. Fluorescence 
was first observed in connection with fluorspar and from this, 
received its name. Fluorescence differs from phosphorescence 
in that the former exists during the continuance of an exciting 
cause while the latter persists after its removal. 

Metallic uranium and all of its compounds are radioactive 
and in fact the phenomena of radioactivity were first discovered 
by Becquerel in 1898 in the uranium compounds. The radio- 
activity of these substances is far less than that of the radium 
compounds and is about the same as that of the thorium salts. 
Again we shall have to postpone a full discussion of the subject 
to a later point. 



Manganese Halogen sub-group 

With the exception of fluorine, which forms no oxy-compounds, 
the members of the halogen family show a variety of valencies 
from one up to seven. Each member of the family forms a 
compound with hydrogen in which the halogen is monovalent. 
They are typical non-metallic elements. 

Manganese is the only member of the family which occupies 
the left-hand column of Group VII in the periodic system. It 
continues the regularities which have been noticed in the pre- 
ceding groups in that it is much more basic than the right-hand 
family; it is unmistakably a metal. It does not form a compound 
with hydrogen analogous to the hydrohalogen acids nor is it 
monovalent in any compound. The strongest resemblance 
between manganese and the halogens is found in the heptoxide, 
Mn 2 O7, and the corresponding permanganic acid, HMnO 4 . 
These are analogous to chlorine heptoxide, C^Oy, and perchloric 
acid, in fact potassium permanganate, KMn0 4 , is isomorphous 
with potassium perchlorate, KC1O 4 . 

Manganese has a rather large number of valencies. It is di-, 
tri-, tetra-, hexa-, and heptavalent and with each alteration of 
valence it changes it chemical relationships, and points of sim- 
ilarity to additional elements become prominent. This makes 
the chemistry of manganese somewhat complicated. It will 
be seen in what follows that in the divalent form it is closely 
related to magnesium; in the trivalent state to chromium and 
iron; in the tetravalent condition to tetravalent lead and tin; 
while the hexavalent manganese is like sulfur in the sulfates and 
the heptavalent resembles chlorine in the perchlorates. 

Occurrence. Manganese is a fairly abundant and widely dis- 
tributed element. Reference to the table on p. 274 will show 
that it stands between barium and strontium in its abundance. 



It is found chiefly in the form of the dioxide, pyrolusite, MnO 2 , 
to which frequent reference has been made in the past since it is 
a much used substance. Other less abundant ores are braunite, 
Mn 2 03, manganite, Mn 2 3 -H 2 O, hausmanite, Mn 3 04, manganese 
blende, MnS, and manganese spar, MnCOs. The metal is ob- 
tained by the reduction of its oxides with carbon in an electric 
furnace or more readily by their reduction with aluminum by the 
Goldschmidt process. Alloys of manganese with iron containing 
from 10 per cent, to 90 per cent, of manganese may be prepared 
in blast furnaces or in electric furnaces. Those containing up 
to 20 per cent, are called spiegeleisen while from 20 per cent, to 
90 per cent, they are known as ferromanganese. These alloys 
are much used in the manufacture of steel. One alloy of man- 
ganese and copper containing about 30 per cent, manganese is 
known as manganese bronze. It possesses great tensile strength 
and is very hard. 

Properties. Pure manganese is a reddish-gray metal which is 
very brittle and is so hard that it will scratch glass. Its melting- 
point is 1,260. It is readily volatilized in the electric furnace. 
It holds its luter in the air but is easily dissolved by acids, even 
by acetic, with the evolution of hydrogen and the formation of 
the divalent or manganous salts. When heated to a little over 
1,200 in a stream of nitrogen it takes fire and burns, forming a 
nitride. In this respect it is much like magnesium. 

Divalent or Manganous Compounds. The manganous salts 
are very much like the salts of magnesium. In many cases the 
two series have about the same solubility and are isomorphous. 
The manganous salts may be prepared by dissolving the metal, 
the carbonate, or any of the oxides in the acid of the salt desired 
and heating. The manganous salts are the only ones which 
are stable in acid solutions and all the others pass into them. 
Nearly all of them have a delicate pink color which is much 
more pronounced in the crystals than in the solutions. The 
hydroxide, Mn(OH) 2 , is thrown down as a white precipitate 
when a soluble hydroxide is added to a manganous salt solution. 
It has about the same solubility and strength as a base as the 
corresponding magnesium hydroxide, and like the latter, it is 
soluble in ammonium salts because of the decrease in the con- 
centration of the hydroxyl through the formation of ammonium 


hydroxide. This solution quickly darkens because of the oxida- 
tion of the manganous ion to the manganic, Mn +++ , by the 
oxygen of the air and the precipitation of manganic hydroxide, 
Mn(OH) 3 , which is too weak a base to be dissolved by ammonium 
salts since it is even weaker than aluminum hydroxide. Man- 
ganous oxide may be prepared by gently heating the hydroxide 
in the absence of air. Manganous carbonate, MnC0 3 , is formed 
as a white precipitate when a soluble carbonate is added to a 
manganous salt solution. It occurs in nature where it is often 
found in isomorphous mixture with magnesium carbonate. The 
precipitated carbonate, like magnesium carbonate, is soluble in 
ammonium salts a salt of the metal, ammonium hydroxide and 
carbonic acid being formed in each case. 

Manganous sulfate, MnS0 4 , usually crystallizes with 4H 2 O, 
but below 6 it has 7H 2 O and is isomorphous with ferrous sulfate, 
while from 7 to 20, the hydrate contains 5H 2 O, and is iso- 
morphous with copper sulfate. It forms the double salt, K 2 S0 4 -- 
MnSO 4 -6H 2 0, with potassium sulfate which is isomorphous with 
the corresponding salt of magnesium and in fact with the whole 
series of these double salts. 

Manganous sulfide, MnS, is the most soluble of the sulfides of 
the heavy metals. Its solubility product is so great that it 
cannot be precipitated by hydrogen sulfide except in the presence 
of hydroxyl or a very large amount of an acetate such as sodium 
acetate. It is usually precipitated by ammonium sulfide and 
generally comes down as an amorphous flesh colored substance 
which is readily soluble in all acids, but it sometimes appears as a 
green powder. When exposed to the air, the sulfide oxidizes 
rapidly and passes back into solution as the sulfate. 

Manganous Ammonium Phosphate, MnNH 4 PO 4 , is formed 
under similar conditions to those for the formation of magnesium 
ammonium phosphate, and has closely the solubility and ap- 
pearance of the latter so that it is not difficult to confuse the two. 

Manganous Borate, MnH 4 (BOs)2, is a salt which is of technical 
importance because it acts as a catalytic agent in the drying of 
oils and varnishes. 

Trivalent Manganese. The trivalent compounds of man- 
ganese or the manganic compounds as they are usually called are 
not as numerous nor as stable as the manganous. A solution of 


the chloride may be obtained by dissolving manganic hydroxide, 
Mn(OH) 3 , is cold hydrochloric acid, but it soon evolves chlorine 
and passes into manganous chloride. The sulfate, Mn 2 (S0 4 ) 3 , 
is a violet-red powder which is very deliquescent and hydrolyzes 
with extreme ease with the precipitation of the hydroxide. It 
forms alums with rubidium and cesium sulfates. A few slightly 
dissociated salts, such as the fluoride, MnF 3 , and the phosphate, 
MnPO 4 , have been prepared which are not hydrolyzed. These 
salts give reddish-violet solutions. Manganese sequioxide, 
Mn 2 03, which is the oxide corresponding to the manganic salts 
occurs in nature as the mineral braunite. 

Tetravalent Manganese. Manganese dioxide is the principal 
tetravalent compound of manganese and is, in fact, the most im- 
portant compound of the element. Like lead dioxide which it 
resembles in a number of ways, it is both feebly basic and acidic. 
When treated with cold hydrochloric acid it is dissolved to form 
the tetrachloride, MnCl 4 , 

Mn0 2 + 4HC1 = MnCl 4 + 2H 2 O 

This hydrolyzes upon dilution with the precipitation of the 
dioxide. When heated, the tetrachloride decomposes into the 
dichloride and chlorine, 

MnCl 4 = MnCl 2 + C1 2 

It was with the aid of these reactions that the chlorine of the 
world was formerly prepared, and large quantities are still 
made in this way although the processes involving the electroly- 
sis of the chlorides have largely replaced it. The tetrahy- 
droxide, Mn(OH) 4 , may be prepared by the action of powerful 
oxidizing agents upon manganous salts in neutral or alkaline 
solutions. In acid solutions the manganese compounds all 
tend to pass into the divalent state, while in alkaline solutions 
they tend to become tetravalent and to be precipitated as the 
dioxide or its hydroxide. This hydroxide is a weak acid and 
forms salts called the manganites of which calcium manganite, 
CaMn 2 O5, may be taken as an example. 

Like lead dioxide, manganese dioxide is a conductor of elec- 
tricity and is used as an oxidizing agent in the ordinary form of 
salammoniac battery known as the Leclanche* cell as well as in 


the slightly modified type called the dry cell. In each of these, 
the oxidizing agent is the dioxide and the reducing agent is 
metallic zinc. The dioxide is reduced to the manganous ion and 
the zinc oxidized to the zinc ion. The electrolyte is ammonium 
chloride solution containing some zinc chloride. The manganese 
dioxide is very slightly soluble, so it is necessary to have a very 
large area for the positive electrode which is of carbon packed 
around with the dioxide. Because of the slight solubility of the 
dioxide, the cell " polarizes' 7 rapidly owing to the exhaustion of 
the dioxide which is in solution and in contact with the carbon 
electrode. It is then necessary to let the cell stand until more 
of the dioxide is dissolved. For this reason, the cell must be 
used on an open circuit. Some idea of the magnitude of the 
dry cell industry may be obtained from the fact that about 50,- 
000,000 of these cells are made in this country per year. 

Hexavalent Manganese, The Manganates. By fusing man- 
ganese dioxide or, in fact, almost any manganese compound with 
potassium hydroxide in contact with the air, or more rapidly 
by the addition of a nitrate or chlorate, a dark green mass con- 
taining potassium manganate, K 2 Mn04, is formed; 

2MnO 2 + 4KOH + 2 = 2K 2 Mn0 4 + 2H 2 O 

This is soluble in water giving a green solution from which 
crystals may be obtained which are isomorphous with potassium 
sulfate. Sodium manganate crystallizes from solution with 
ten molecules of water, Na 2 MnO 4 -10H 2 O, and is isomorphous 
with the decahydrate of sodium sulfate. From these two ex- 
amples, it will be seen that the manganates are analogous to 
the sulfates, and consequently hexavalent manganese to hexa- 
valent sulfur. Solutions of manganates have an exceedingly 
intense green color, but are stable only in the presence of a 
hydroxide. Manganic acid cannot be prepared because it at 
once decomposes into permanganic acid and the dioxide, but its 
anhydride, MnO 3 , has been obtained as a reddish deliquescent 
powder which is very unstable. 

Heptavalent Manganese, The Permanganates. When a solu- 
tion of a manganate is diluted or acidulated, the color changes 
through blue and violet to purple and at the same time man- 
ganese dioxide is precipitated. The change is due to the oxida- 


tion of a portion of the manganate to the permanganate at the 
expense of the remainder which is reduced to the dioxide, 

3K 2 Mn0 4 + 2H 2 = 2KMn0 4 + 4KOH -f Mn0 2 

This reaction can be easily understood if the equation be written 
in the ionic form omitting the ions which do not enter directly 
into the reaction, 

3Mn0 4 + 2H+ = 2MnOr + 20ET + Mn0 

Since H + is consumed in the reaction and OH~ produced, this 
reaction will take place much more readily when the solution is 
acidulated. Even a very weak acid will bring about this change, 
so a solution of a manganate gradually becomes purple upon 
exposure to the air from the absorption of carbon dioxide and 
the formation of carbonic acid. 

The transformation of the manganate into the permanganate 
is oxidation and may be brought about by passing chlorine into 
the solution of the manganate. In this case, the chlorine is 
reduced to the ion Cl" and the whole of the manganate is oxi- 
dized to the permanganate so that manganese dioxide is not 

2K 2 MnO 4 + C1 2 = 2KMnO 4 + 2KC1 

In the permanganate ion Mn0 4 ", since each oxygen carries 
two negative charges, see p. 184, the manganese must have seven 
positive charges while in the manganate, MnO 4 , on the same 
basis the manganese carries six, and in the dioxide, Mn0 2 , four 
positive charges. Then when one gram ion of the manganate is 
reduced to the dioxide the manganese gives up two + charges in 
passing from + 6 to -f- 4 and these two + charges are just 
sufficient to raise the manganese in two gram ions of the man- 
ganate from + 6 to +. 7 or in other words, to oxidize them to the 
permanganate. Reference to the equations given above will 
show that this is the case, and three gram moles of manganate 
yield two gram moles of permanganate to one of manganese 
dioxide. Each atomic weight of chlorine gives up a -f charge 
in passing into the chlorine ion Cl", and hence a mole of chlorine 
can oxidize two moles of manganate. 

The permanganates are powerful oxidizing agents and disin- 


fectants and are manufactured in large quantities for such pur- 
poses. The sodium salt is the cheapest if a pure article is not 
demanded. It is so very soluble that it is not easily obtained 
pure. For laboratory purposes the potassium salt is generally 
used. This crystallizes well in purple crystals with a greenish 
luster, and is isomorphous with potassium perchlorate. 

Solutions of potassium permanganate of known concentration 
are much used in the laboratory in the volumetric determination 
of easily oxidizable substances such as ferrous iron, oxalic acid or 
manganous salts, since the permanganate acts in a rapid and 
definite manner with these substances. In carrying out the 
operation, the permanganate solution is added from a burette to 
the solution to be tested and the volume necessary to give to the 
latter the purple color of the permanganate is noted. The opera- 
tion is especially convenient since, owing to the strong color of 
the permanganate, no other indicator is necessary. When the 
permanganate is used in this way, it is reduced to the manganous 
ion in acid solutions or to manganese dioxide in neutral or alka- 
line solutions. The titration of ferrous iron or oxalic acid is 
carried out in acid solutions, so the manganese in passing from 
the heptavalent to the divalent state gives up five + charges, 
and hence one mole of permanganate is able to oxidize five moles 
of a reducing agent which takes up one -f- charge per mole, or 
two moles of the permanganate can oxidize five moles of reducing 
agent which takes up two + charges as is the case with the 
oxalate ion. 

The following equations represent the reaction mentioned 

5Fe+ + + MnO 4 ~ + 8H+ = 5Fe++++ Mn++ + 4H 2 O 

10FeS0 4 + 2KMn0 4 + 8H 2 S0 4 = 5Fe 2 (SO 4 ) 3 + K 2 S0 4 + 

2MnS0 4 + 8H 2 

5C 2 O 4 --+ 2MnOr + 16H+ = 10C0 2 + 2Mn++ + 8H 2 O 

5H 2 C 2 4 + 2KMn0 4 + 6HC1 = 10CO 2 + 2KC1 + 
2MnCl 2 + 8H 2 

As has been mentioned, the permanganate is reduced to 
manganese dioxide in neutral or alkaline solutions, and also in 


these same solutions, the manganous ion tends to oxidize to the 
tetravalent condition and to be precipitated as the dioxide. The 
result of these tendencies is that manganous compounds may be 
titrated by permanganate in neutral solutions. In going from 
the permanganate to the dioxide, the manganese gives up three 
+ charges, and the manganous in going to the dioxide takes up 
two + charges; hence two moles of permanganate will oxidize 
three moles of manganous salt as shown in the following equation, 

2KMn0 4 + 3MnS0 4 + 2H 2 = K 2 S0 4 + 2H 2 S0 4 + 5Mn0 2 

The solution tends to become acid which would interfere with 
the reaction, so the titration is carried on in the presence of an 
excess of zinc oxide which neutralizes the acid as it is formed. 
The behavior of permanganate as an oxidizing agent can be 
readily kept in mind by remembering that in acid solutions it 
gives up five plus charges and goes to the manganous state, 
while in neutral or alkaline solutions it gives up three charges 
and becomes manganese dioxide. 

When dry potassium permanganate is heated to 240 it de- 
composes into oxygen the manganate and manganese dioxide, 

2KMnO 4 = 2 + K 2 Mn0 4 + Mn0 2 

This gives purer oxygen than the decomposition of the chlorate 
which generally yields some chlorine. 

A solution of permanganic acid may be prepared by the second 
general method for the preparation of acids, using barium per- 
manganate and sulfuric acid. It has the deep purple color of 
the permanganates and decomposes upon boiling or exposure 
to the light. An impure solution of the acid is also formed when- 
ever a manganese compound is boiled with moderately dilute 
sulfuric or nitric acid and red lead or lead dioxide. The equation 
in the latter case is, 

2MnSO 4 + 5Pb0 2 + 3H 2 SO 4 = 2HMnO 4 + 5PbS0 4 + 2H 2 O 

Because of the strong and very characteristic color of the per- 
manganate ion, this constitutes one of the best tests for man- 
ganese. It is necessary, however, to have the solution highly 
acid and to use a very small quantity of manganese as otherwise 
the action stops with the formation of the dioxide. Manganese 


heptoxide, Mr^O?, is formed by dissolving potassium perman- 
ganate in cooled concentrated sulfuric acid and carefully adding 
water when the oxide separates as a dark reddish-brown liquid 
which decomposes with the evolution of oxygen and explodes at 
the slightest provocation. 

Analytical Properties of Manganese. The solubility of man- 
ganese sulfide in acids and its precipitation by ammonium sul- 
fide puts this element in an analytical group with zinc, cobalt, 
and nickel. From these it is distinguished by the pink color of 
its sulfide, by the formation of permanganic acid as described 
just above, by the green mass of manganate which is formed by 
fusion with an alkali and a nitrate, or by the amethyst color 
which all manganese compounds give in a borax bead heated in 
the oxidizing flame. 


Fe Co Ni 
Ru Rh Pd 
Os Ir Pt 

An examination of Group VIII of the periodic system will 
show that it differs radically from the other groups in that it is 
made up of three families of closely related metals in which 
the similar elements occur side by side, and have about the same 
atomic weight, instead of forming a perpendicular column 
with widely different atomic weights. There is, however, some 
similarity in the perpendicular columns, particularly between 
the last two families. 

The three sub-families of the group are: 

The iron sub-group: iron, atomic weight 55.84, cobalt, atomic 
weight, 58.97, and nickel, atomic weight, 58.68. 

The ruthenium sub-group: ruthenium, atomic weight 101.7, 
rhodium, atomic weight 102.9, and palladium, atomic weight 

The platinum sub-group: osmium, atomic weight 190.9, irid- 
ium, atomic weight 193.1, and platinum, atomic weight 195.2. 

All the members of the group are true metals and are some- 
what remarkable for the number of complex compounds which 
they form. 

From the gradual increase in the maximum valence which 
has been observed in the preceding groups, it might be antici- 
pated that the valence of the members of this group would reach 
eight. The only known compounds of this kind are the tet- 
roxides of ruthenium and osmium, RuO4 and Os04 and osmium 
octofluoride, OsF 8 . 


Iron forms three series of compounds: the ferrates in which 
it is hexavalent and resembles manganese, and chromium in the 
so 466 


manganates and chromates; the ferric salts in which it is triva- 
lent and is like aluminum, and trivalent chromium or mangan- 
ese; and the ferrous salts in which it is divalent and resembles 
the metals of the magnesium groups as well as cupric, chro- 
mous, and manganous compounds. Cobalt forms two series of 
salts similar to the ferric and ferrous, but the cobaltous are, in 
general, the more stable. Nickel forms only one series of salts, 
the nickelous, but shows some tendencies toward higher valence 
in the oxide, Ni 2 0s. The members of this family are all mag- 
netic, have high melting-points and decompose steam at a high 


General. This is truly the age of iron, as may be seen from 
the fact that the United States alone produces yearly about 30,- 
000,000 long tons of the metal. A moment's thought will show 
that our present civilization would be very different if we did 
not have iron and steel in such abundance. Iron has been 
known and used from prehistoric times. Many meteors are 
composed chiefly of metallic iron, and the earliest iron tools and 
weapons were probably made from this source. The ores of 
iron are easily reduced by carbon, and even very primitive 
peoples have known how to do this in a small way, but the 
production on anything like the modern scale has only been 
possible since the invention of the blast furnace. 

Occurrence. Native iron is found in meteorites, and in certain 
deposits in Greenland; but all that is used is obtained from 
compounds. The principal ores are hematite, ferric oxide, 
Fe 2 0s; brown iron ore or hydrated ferric oxide, Fe 2 O3nH 2 O, 
this is also called brown hematite, bog ore or limonite; magnetite, 
magnetic iron oxide, FesO^ and siderite, spathic iron or kidney 
ore, ferrous carbonate, FeCOa. Hematite furnishes about nine- 
tenths of the ore. Another iron mineral is pyrite, FeS 2 ; this 
is mined in large quantities and burned as a source of sulfur 
dioxide in the manufacture of sulfuric acid. The residue con- 
sists principally of Fe 2 03, but is too finely powdered and con- 
tains too much sulfur to use as an ore of iron. It has recently 
been found that this material may be sintered in a rotary* kiln 
similar to that used in the manufacture of Portland cement, 


and brought into a suitable mechanical state; at the same 
time the sulfur content is reduced to a point where it is no longer 
objectionable. This makes what was a waste product, a valuable 
source of iron. More than 50,000,000 tons of iron ore are mined 
per year in this country, the greater part of this comes from a 
region around the south and west of Lake Superior. It is mined 
in great open pits with steam shovels, loaded directly into cars 
and taken to the harbors on the lake, transferred to ships and 
carried to the lower lakes, where it is either unloaded and smelted 
on- the spot, or taken by train to the great iron smelters around 
Pittsburg which are conveniently located with respect to the 
coal fields and manufacturing centers. Within recent times a 
number of mammoth iron and steel plants have been located on 
the shores of the Great Lakes, so it looks as though Pittsburg 
may lose its pre-eminence in this industry. Another important 
iron-producing region is around Birmingham, Alabama, which 
has near at hand both iron ore and coal. Everything must be 
done as cheaply as possible because the average price of pig iron 
in normal times is about $15 per long ton. 

Metallurgy of Iron. The manufacture of all forms of iron and 
steel begins with the making of cast iron in a blast furnace. 
This depends for its usefulness upon the fact that by properly 
selecting and proportioning the materials put into the furnace, 
the impurities in the ore and the products of the reactions may 
all be removed as gases or liquids, and the operation made con- 
tinuous for a long time. The materials used are coke, iron ore, 
and limestone. The coke burns and gives the necessary high 
temperature, and also furnishes the requisite reducing agent, 
carbon monoxide. The limestone is changed to lime which 
reacts with the silicious impurities in the ore to form an easily 
fusible slag, consisting of calcium and aluminum silicates, which 
is removed from the bottom of the furnace in the liquid state. 

A modern blast furnace is a very large affair; 80 to 100 ft. in 
height by about 22 ft. in greatest diameter. Its shape is approxi- 
mately that of two truncated cones set base to base. It is made 
of steel plates and lined with very infusible silicious fire brick. 

A diagrammatic sketch of a blast furnace is shown in Fig. 70. 
The limestone, ore and coke in the proper proportions are charged 
in at the top. A blast of hot dry air is blown in through 8 to 16 



tuyeres near the bottom. The coke nearest the tuyeres burns 
to carbon dioxide, which is then reduced by that farther away to 
carbon monoxide, which passing up through the iron oxide 
reduces it first from Fe 2 O 3 to Fe 3 O 4 , then to FeO and finally 
to the metal. The latter then takes up enough carbon to lower 
its melting-point from 1,505, that of pure iron, to 1,100-1,200, 
that of cast iron, and melts and collects in the lower part of the 

FIG. 70. 

cylindrical portion of the furnace which is known as the hearth. 
The slag collects in the fused state just above the iron and is 
tapped off once in a while. The iron is also removed from the 
furnace through a tap hole at intervals, and either cast into bars 
called pigs, or kept in the molten state in large mixers until it 
is needed for the conversion into steel in Bessemer converters 
or open hearth furnaces. Fig. 71 shows the substances and 
their weights used to produce a long ton of pig iron; the equa- 



tions for the reactions which take place; where they occur; the 
temperature at the various parts of the furnace; and the dis- 
posal of the products. It should, therefore, receive careful 
study. The gases that escape from the top of the furnace con- 
tain under the best conditions approximately 16 per cent. COz, 
24 per cent. CO and 60 per cent. N 2 . The carbon monoxide is 

Pig Iron I 

l->Mineral Wool 

\-*~ Castings 
Wrought Iron 

FIG. 71. 

combustible and the gas produced in a furnace represents thou- 
sands of horse-power. About one-third of it goes to the stoves 
for heating the blast; the rest is used to develop the power 
needed around the plant being burned under boilers or better in 
gas engines. Heating the blast saves a great deal of fuel; another 
great saving is made by drying the air, before it is heated, by 
cooling it to a low temperature with an ammonia refrigerating 


plant, thus freezing out the water. This helps, because the water 
vapor in the air would react with the. coke with the formation of 
carbon monoxide and hydrogen and the absorption of much 
heat. The saving in fuel caused by drying amounts to 20 per 
cent, and the output of the blast furnace is increased by the 
same amount. 

The cast iron as it comes from the blast furnace is impure and 
contains from 4 to 11 per cent, of foreign substances, chiefly 
carbon, silicon, manganese, sulfur and phosphorus. These sub- 
stances lower the melting-point of the iron and make it crystal- 
line and brittle over all ranges of temperature. Cast iron, there- 
fore, is not suitable for any purpose which requires considerable 
tensile strength, but may be used for castings which will be sub- 
jected to compression without much impact. There are two 
varieties of cast iron white and gray. The white kind is made 
by rapid cooling; it contains the carbon in chemical combina- 
tion and is very hard and brittle. The gray iron is made by 
slow cooling. It is tougher and softer than the white and con- 
tains a large part of the carbon in the form of crystals of graphite 
scattered through the mass. 

By removing these impurities more or less completely, the 
iron may be converted into steel or wrought iron and thereby 
acquires much more valuable properties. More than three- 
quarters of the pig iron manufactured is subjected to such puri- 
fication processes. The first step in each of these is to take ad- 
vantage of the fact that the impurities mentioned are more easily 
oxidized than the iron, and hence will be eliminated before the 
latter is attacked in case the crude substance is exposed to oxidation. 
This oxidation is done either by the oxygen of the air or by 
that from iron oxide. It is difficult to draw the line between 
wrought iron and steel because the division is arbitrary, but the 
following is in common usage. Wrought iron is slag-bearing, 
contains less than 0.12 per cent, carbon, and does not harden 
materially when suddenly cooled. It is made in puddling 
furnaces. Steel is iron which is malleable at least over some 
range of temperature, contains from 0.12 to 2.2 per cent, of car- 
bon, and which when suddenly cooled, is materially hardened. 
It is made in Bessemer converters, open hearth furnaces, or in 



Bessemer Process. Kelly, an American, invented and pat- 
ented in 1852 the process of purifying molten pig iron by blow- 
ing air through it. Bessemer, an Englishman, patented in 1855 
the converter or vessel best adapted to carry out the process. 
After litigation, the Kelly interests sold out to the Bessemer, and 
the process now bears the name of the latter. 

FIG. 72. 

The converter used is a pear-shaped vessel, built of boiler plate 
and lined with silicious material, so mounted on trunions that it 
may be tipped through a very large angle, and provided with 
about 250 small holes in the bottom through which air may be 
blown (see Fig. 72). To start the blow, the converter is turned 
on its side and ten to fifteen tons of molten pig iron poured in. 


The air is turned on and the converter straightened up so that 
the air passes through the molten iron, buring first the silicon 
to Si02, then the manganese to MnO, and lastly the carbon to 
CO. This burns at the mouth of the converter forming a great 
flame which gradually dies down and disappears at the end of 
the blow. The whole process is completed in from six to ten 
minutes. The heat produced raises the temperature some 300 
which is favorable to the success of the operation, because 
removing the impurities raises the melting-point of the iron. By 
the time that the carbon is burned out, much of the iron is oxi- 
dized to FeO. This makes the product brittle and unfit for use. 
This fault is overcome by adding at the end of the blow such a 
quantity of spiegeleisen (an alloy of iron, manganese and car- 
bon) as will furnish the manganese and carbon necessary to 
reduce the ferrous oxide and to bring the manganese, silicon, 
and carbon content of the whole mass to that required in the 
grades of steel desired. 

This process using the acid or silicious lining leaves in the 
steel all the phosphorus and sulfur present in the pig iron. But 
by using basic linings made of lime or calcined dolomite, CaCO 3 - 
MgCOa, and adding lime with the charge, it is possible to remove 
the phosphorus and a part of the sulfur as calcium or magnesium 
phosphate or sulfide. The basic lining is less durable than the 
acid and the process more expensive, so it has been abandoned 
in America, although the properties of steel containing phos- 
phorus or sulfur are so objectionable that only the purest 
and . most expensive pig iron can be refined in acid lined 

After the addition of the spiegeleisen, the steel is poured 
out of the converter into a ladle and then tapped into cast 
iron moulds where it solidifies, forming ingots which are 
reheated and rolled into the shapes desired rails, beams, 
rods, etc. 

A very superior quality of steel can be produced by adding a 
small quantity of an alloy of iron and titanium soon after the 
addition of the spiegeleisen. The titanium completes the 
reduction of the oxides, combines with any nitrogen present, 
partially removes the phosphorus and sulfur, and greatly im- 
proves the strength, malleability, and toughness of the steel. 



Open Hearth or Siemens-Martin Process. The open hearth 
process for steel making is even more important than the Besse- 
mer as may be seen from the fact that this country produces 
yearly 16,500,000 tons by the former as against 9,400,000 tons 
by the latter. This is due largely to the fact that a basic lining 
may be economically used in an open hearth furnace, and hence 
steel may be cheaply made from pig iron which contains too 
much phosphorus for the acid Bessemer converter. 

The furnace used (see Fig. 73) is of the type known as regenera- 
tive because the air and the producer gas which is used as fuel 
are preheated by the waste heat of the escaping products of com- 
bustion, with the result that the temperature reached is very 

Floor Line 

FIG. 73. 

To secure these results, the air and gas are blown in separately 
through chambers filled with checkerwork of hot brick. They 
meet and burn just over the bed of the furnace containing 
the materials for forming the steel. The products of com- 
bustion pass out at the other side of the furnace through chambers 
which are duplicates of the ones through which the gases entered. 
In this way the checkerwork of brick in the exit chamber is 
intensely heated. After a time the direction of the passage of 
the gases through the furnace is reversed, and the brick in what 
had been the exit chambers now warms the incoming air and 
gas. This reversal is continued at regular intervals. Because 
of this preheating, the temperature reached in the furnace is 
very high. This type of furnace was invented by Siemens 
Brothers and applied to steel-making by Martin; hence the name 
Siemens-Martin process. 


The materials used are pig iron, soft scrap steel, iron ore and 
in the basic process, lime. The purification of the pig iron is 
effected partly by an excess of air in the flame which plays over 
the charge, but more by the oxidizing action of the iron oxide 
in the ore. When the oxide is reduced, it of course increases 
the amount of iron present. The scrap steel is added to dilute 
the impurities and so hasten the process. The lime makes a 
very basic slag in the presence of which calcium phosphate 
and calcium sulfide are formed with the elimination of phos- 
phorus and sulfur. Lime can be used only when the lining is 
made of calcium and magnesium oxide or the latter alone. The 
process is much slower than the Bessemer, requiring from six 
to ten hours, but the capacity of the furnace is large, usually 
50 to 75 tons. Because of its slowness the changes can be 
followed by physical and chemical tests, and hence a more uni- 
form and higher grade product can be produced than by the 
Bessemer converter. After the purification, the charge is tapped 
out into a ladle, slag is removed as completely as possible, 
and the iron recarburized to the desired degree by the addition of 
ferromanganese (an iron alloy rich in manganese) and anthracite 
coal or charcoal. It is then cast into ingots and worked up as in 
the other process. 

The acid process is used to make about 1,200,000 tons of 
the 16,500,000 tons produced in the open hearth. It differs 
mainly in using a silicious lining and omitting the lime from the 
charge. The materials used must be very low in phosphorus 
and sulfur. 

Crucible Steel. By melting wrought iron with a proper 
amount of charcoal or cast iron in large clay or graphite crucibles, 
a very fine grade of high carbon steel is made which is especially 
desirable for making tools, knives, springs, and other special uses. 
The steel made by this method usually has a carbon content of 
from 0.75 to 1.50 per cent, and is called crucible steel. 

Electric furnace steel is practically the same as crucible steel. 
The main difference being that melting is carried out in an 
electric furnace instead of a crucible, and that larger charges 
can be worked at a lower cost. Because of the high temperature 
obtainable in these, they are often used to give Bessemer or 
open-hearth steel a super-refining which removes phosphorus 


and sulfur, and greatly improves the quality. The electricity 
simply serves as a source of heat and there is no electrolysis in 
the process. 

Within recent years, it has been found that the addition of one 
or more of the following metals to steel will increase its valuable 
qualities, or impart new ones; nickel, manganese, copper, chro- 
mium, vanadium, molybdenum, tungsten, titanium, and silicon; 
such steels are called alloy steels, and are extensively used. 
Manganese steel is very tough. Nickel increases the tenacity 
and elastic limit. Chromium increases the hardness. 

Wrought Iron. A third method of purifying pig iron is the 
puddling process, invented by Henry Cort in 1784. This 
consists in melting pig iron in small lots of 500 to 1,500 Ib. in a 
reverberatory furnace heated by coal or gas. 

The hearth of the furnace is lined with roll scale or other 
material rich in iron oxide which with the excess oxygen in the 
furnace gases, serves to oxidize silicon, manganese, carbon, phos- 
phorus and sulfur. The silica, manganous oxide, and phos- 
phorous pentoxide produced, unite with some iron oxide formed 
by oxidation of iron, making a very basic slag, and effectually 
removing them from the iron. The melting-point of the iron 
which is about 1,200 in the impure state at the beginning, gradu- 
ally rises as the impurities are eliminated and finally becomes 
so high that the pure iron cannot remain liquid at the highest- 
temperature of the furnace. It separates, therefore, as a spongy 
plastic mass, honey-combed with slag. This is gathered together 
and worked into balls weighing about 125 Ib., which are then 
removed from the furnace and passed through a mechanical 
squeezer to expel slag. The iron is rolled into rough bars, which 
are cut up, piled together, reheated to welding temperatures, 
and rolled into bars and other shapes. 

The product of the puddling furnace, called wrought iron, is 
very nearly free from impurities, other than slag, and is the purest 
commercial iron. 

It is very soft and malleable, welds easily, and is characterized 
by a fibrous structure, which easily distinguishes it from low 
carbon steel, which is rolled into form from cast ingots nearly 
free from slag. This structure is due to particles of slag enclosing 
those of iron and causing the iron particles to stretch separately 


during the rolling process. The fibrous structure, it is thought, 
increases the strength of the iron. 

Physical Properties of Iron. Pure iron is the most magnetic 
of the metals. It has a density of 7.8, is silvery white, takes a 
high polish and is very ductile and malleable. It may be welded 
at a bright red heat and melts at 1,505. 

Pure iron is soft and cannot be hardened, but steel can be 
hardened by heating to a high temperature and suddenly cool- 
ing in water. If this hardened steel is reheated, it may be 
softened or tempered as it is called. The explanation is found in 
the fact that iron forms a carbide, Fe 3 C, called cementite which 
is soluble to a limited extent in iron at temperatures above 725 
forming solid solutions. The solubility increases with the 
temperature reaching 2.2 per cent, carbon at 1,145. If these 
solid solutions are quickly cooled by plunging the steel into 
water, the cementite remains in solution, and the resulting 
solid solution is hard. The more cementite in the solution the 
harder the steel is. If this hardened steel be reheated, the 
cementite begins to separate and the steel gets softer. By 
controlling the temperature and the duration of reheating, the 
steel may be brought to any desired degree of hardness between 
that of the chilled steel and that of pure iron. The softening 
process may be stopped at any point by suddenly cooling the 
steel. Experienced workmen judge the hardness by the colors 
of the film of oxide on the steel which are a rough measure of 
the temperature to which the iron is heated. 

Chemical Properties of Iron. Iron is stable in dry but is 
attacked by moist air and gradually becomes covered with iron 
rust, 2Fe20 3 -3H 2 O, which does not protect the underlying 
metal from further action. Something of a discussion of the 
mechanism of the rusting of iron will be given at the close of the 
treatment of this metal. When heated in the air or steam, 
iron becomes converted into the magnetic oxide, Fe 3 4 , which 
when formed under certain conditions clings tightly to the sur- 
face of the metal and protects it from further oxidation (Baruff s 
process for the prevention of rust). Dilute hydrochloric or 
sulfuric acids dissolve iron with the evolution of hydrogen and 
the formation of ferrous salts: the hydrogen usually contains 
some hydrocarbons as well as compounds of phosphorus and 


sulfur and has a disagreeable odor. When the iron dissolves 
in cold dilute nitric acid, ferrous nitrate and ammonium nitrate 
are formed; while with somewhat more concentrated acid, ferric 
nitrate and oxides of nitrogen are produced. When iron is dipped 
into very concentrated nitric acid, it almost instantly becomes 
passive, that is passes into a condition in which it is not attacked 
by nitric acid, dilute or concentrated, and does not precipitate 
copper from cupric solutions. In fact, it acts as though it 
were moved in the potential series to a position near that of 
platinum. Potassium dichromate, chloric, bromic, iodic acids 
and other powerful oxidizing agents will also induce the passivity 
of iron. There is no visible change in the iron. The passivity 
is destroyed by contact with reducing agents, scratching the 
surface or even by simply placing the iron in a very strong mag- 
netic field. An entirely consistent and satisfactory explanation 
of this phenomenon is lacking. 

Ferrous Compounds. The ferrous compounds are very much 
like those of magnesium and especially like manganous salts 
with the exception that the trivalent compounds of iron are more 
stable than the divalent either in acid or alkaline solution, while 
the reverse is true for manganese in acid solution. The ferrous 
salts even in acid solutions are oxidized in the air to ferric; in 
alkaline solutions, the change takes place more readily. When 
a perfectly pure ferrous salt solution is acted upon by a soluble 
base in the entire absence of oxygen, a white precipitate of 
ferrous hydroxide, Fe(OH) 2 , is obtained. On exposure to the 
air, it becomes first dirty green and finally brown from oxida- 
tion to ferric hydroxide. Ferrous hydroxide has about the solu- 
bility and strength of magnesium and manganous hydroxides and 
hence, like these is soluble in an excess of ammonium salts and 
for the same reason. It is not soluble in excess of the alkalies 
or of ammonium hydroxide, but is, of course, in acids. From 
its solubility in ammonium salts, it follows that it is a fairly 
strong base and its salts are but little hydrolyzed, so those of even 
as weak an acid as carbonic can be easily prepared. Ferrous ox- 
ide, FeO, corresponding to the hydroxide is most easily made by 
heating the oxalate in the absence of air. It has a black color. 

Ferrous Sulfate. Ferrous sulfate, FeSO 4 -7H 2 O, is the best 
known and most largely used of the ferrous salts. In its crude 


form it is often called green vitriol or copperas. It is made on a 
very large scale by the oxidation of moist pyrite. 

2FeS 2 + 70 2 + 2H 2 = 2FeSO 4 + 2H 2 SO 4 

The ferrous sulfate and sulfuric acids which are formed are 
leached out, the acid is allowed to act on scrap iron, and the solu- 
tion concentrated to crystallization. Rather large quantities of 
ferrous sulfate are obtained as a byproduct in plants where there 
is occasion to clean or "pickle" the surface of much iron by the 
action of sulfuric acid. 

The sulfate is used as a deodorant and disinfectant, in the 
manufacture of dyes, in tanning, in the making of inks, and in 
the clarification of water. 

Ferrous Sulfide. Ferrous sulfide, FeS, is formed by the direct 
union of the elements. This takes place very readily when the 
two are heated together. The sulfide has a slight metallic luster. 
It is readily dissolved by hydrochloric or sulfuric acid with the 
formation of hydrogen sulfide and the corresponding ferrous salt. 
The usefulness of hydrogen sulfide as a reagent makes ferrous 
sulfide important also. It is precipitated by ammonium sulfide 
as a black powder easily soluble in acids, but is not precipitated 
by hydrogen sulfide from acid solutions. 

Other Ferrous Compounds. Ferrous chloride, FeCU, is much 
more soluble than ferrous sulfate, and on this account it is advis- 
able to use hydrochloric acid instead of sulfuric in the generation 
of hydrogen sulfide, as the apparatus is not so likely to become 
clogged with crystals of the salt. There is nothing especially 
noteworthy concerning the bromide and the iodide except that 
ferrous iodide is the only stable iodide of iron, the ferric salt 
decomposing into the ferrous and free iodine. Ferrous carbonate 
occurs in nature as spathic iron ore. As it is obtained in the 
laboratory, it is partially hydrolyzed and is easily oxidized to 
the ferric hydroxide. It is more soluble in water containing 
carbon dioxide in solution than in pure water. 

Ferric Compounds. As has been mentioned, the ferric com- 
pounds are formed from the ferrous by oxidation. This takes 
place more readily in alkaline than in acid solutions, but the 
oxygen of the air is able to bring about the change in either 


Any ferric salt in solution will be reduced to the ferrous 
state by the stronger reducing metals such as zinc, magnesium, 
iron, etc. ; or by hydrogen sulfide, stannous chloride, or iodine as 
ion. Some of these properties are made use of in the quantitative 
determination of iron, it being reduced to the ferrous state and 
then titrated with permanganate. 

Ferric Hydroxide. Solutions of bases when added to a solution 
of a ferric salt, produce a brown precipitate of ferric hydroxide, 
Fe(OH) 3 . This precipitate is very slimy when formed in the 
cold, but becomes more compact upon boiling. It has a marked 
tendency to pass into colloidal solution especially in the presence 
of ferric chloride. By separating such a solution from pure 
water by parchment paper, dialysis will take place and the ferric 
chloride will gradually diffuse into the water leaving a pure solu- 
tion of colloidal ferric hydroxide. This is called dialyzed iron 
and has a strong red color. Like other colloids it has practically 
no effect upon the boiling-point of the water, and is precipitated 
by many neutral salts. It is a positive colloid and di- and trivalent 
anions are much more effective than monovalent ones. Ferric 
hydroxide is soluble in acids, but not in excess of dilute alkalies; 
or in ammonia or ammonium salts. It is a very weak base, and 
its salts are as highly hydrolyzed as the salts of^aluminum; 
hence, their solutions are decidedly acid % The ferric ion appears 
to be colorless, but the ferric salt solutions are usually colored, 
presumably, in most cases, from the hydroxide formed by 

When the hydroxide is dried and heated, it passes into ferric 
oxide, Fe20 3 , which in finely divided condition is known as 
" rouge" and "Venetian red/' and is used as a polishing material 
and pigment. Although ferric hydroxide will not dissolve in 
dilute alkalies it is somewhat soluble in very concentrated solu- 
tions of sodium or potassium hydroxide, forming unstable fer- 
rites. The dark colored substance formed during the early stages 
of the oxidation of ferrous hydroxide is probably a hydrated fer- 
rous ferrite; and magnetic iron oxide, Fe 3 04, is in all probability 
the anhydrous compound and should be written Fe(Fe02)2- 
This is a valuable iron ore and remarkable because of its being 
attracted by magnets. Some specimens act as magnets and are 
then known as lodestones. Calcium, magnesium, and zinc fer- 


rites of the general formula M(Fe0 2 )2, are known. These like 
the ferrous ferrite are magnetic. 

Ferric Chloride. When chlorine is passed over heated iron, 
anhydrous ferric chloride, FeCl 3 , sublimes and crystallizes in 
dark green scales which have a metallic luster and are red by 
transmitted light. It dissolves in water with the evolution 
of heat and from the solution, a number of hydrates may be 
obtained, the most common one is FeCl 3 -6H 2 O. This is most 
readily prepared, by oxidizing ferrous chloride with chlorine, 
evaporating the solution until it has the composition of the hexa- 
hydrate and allowing it to cool. The hydrate melts at 37 
and is very soluble and deliquescent. Solutions of ferric chloride 
are considerably hydrolyzed, more so at high temperatures then 
at low, and react strongly acid. The chloride is soluble in alcohol 
and ether as well as in water. It is used in medicine. 

Ferric Sulfate. Ferric sulfate, Fe 2 (S0 4 ) 3> may be prepared 
as a yellowish white solid by adding the proper amount of sulf uric 
acid to ferrous sulfate in solution and then oxidizing with nitric 

6FeSO 4 + 3H 2 SO 4 + 2HN0 3 = 3Fe 2 (S0 4 ) 3 + 2NO + 4H 2 O 

and evaporating to dry ness. With potassium and ammonium 
sulfates, it forms alums f such as ammonium iron alum, NH 4 Fe- 
(SO 4 ) 2 -12H 2 O. When pure these are almost colorless, but 
generally have a violet tint due perhaps to the presence of a trace 
of manganese. Basic ferric sulfates are known and are used in 
medicine to stop the flow of blood from wounds. 

Ferric Thiocyanate. Ferric thiocyanate, Fe(NCS) 3, is a slightly 
dissociated salt which is formed by the interaction of a ferric salt 
and a soluble thiocyanate. It has a most intense blood red color 
which is visible even at high dilutions. The appearance of this 
color upon the addition of potassium thiocyanate is a very 
delicate test for ferric iron. Pure ferrous compounds give no 

Ferric Sulfide. Ferric sulfide, Fe2S 3 , is thrown down as a 
black precipitate when a ferric salt is added to an excess of 
ammonium sulfide, but when the ferric salt is in excess, reduc- 
tion takes place with the formation of ferrous sulfide and sulfur. 
Ferric sulfide is also formed by the direct union of the elements. 


Pyrite. Pyrite or " Fool's gold," FeS 2 , is the most stable of 
the sulfides of iron at ordinary temperatures. It is found in 
very large quantities in nature and is an important source of the 
sulfur for sulfuric acid. It occurs in golden colored crystals 
which appear in a great variety of forms belonging to the regular 
systems, cubes and octahedrons being common. 

Ferric Phoshate. Ferric pho^hate, FeP0 4 , is so slightly soluble 
that it is precipitated upon the addition of sodium phosphate to 
a ferric solution even in the presence of acetic acid in which most 
phosphates are soluble. Advantage is sometimes taken of this 
property to remove the phosphate ion from solutions. It is a 
white slimy precipitate soluble in strong acids. 

Cyanogen Compounds. When potassium cyanide is added to 
a ferrous solution a yellow precipitate is formed which is ferrous 
cyanide. It dissolves in an excess of potassium cyanide, form- 
ing the typical complex salt, potassium ferrocyanide or yellow 
prussiate of potash, K4Fe(NC) 6 . This is easily soluble in 
water and the solution shows the properties of other potassium 
salts, but not those of either the ferrous salts or the cyanides; 
on the contrary, it has an entirely different set of properties. 
The salt for example is nothing like as poisonous as potassium 
cyanide. This is expressed by saying that it dissociates into potas- 
sium as ion and the complex ferrocyanogen ion, Fe(NC) 6 . 

This ferrocyanogen ion is formed whenever the ferrous ion is 
boiled with the cyanogen in alkaline solution. The potassium 
salt has the composition, K 4 Fe(NC) 6 -3H 2 O, and the sodium, 
Na4Fe(NC) 6 -10H2O. Many of the ferrocyanides are insoluble 
and may be prepared by precipitation. Ferric ferrocyanide, 
Fe4(Fe(NC) 6 )3, has a dark blue color and is known as " Prussian 
blue." The formation of this dark blue precipitate is a very 
delicate test for ferric iron and it is indeed remarkable that a 
solution containing iron can be used as a reagent for iron. When 
potassium ferrocyanide is added to a ferrous salt solution in the 
absence of oxygen, a white precipitate having the composition, 
K2FeFe(NC)e, is formed which quickly oxidizes and becomes 

Ferricyanides. Ferricyanides may be easily formed from the 
ferrocyanides by the action of oxidizing agents. The potassium 



salt, KsFe(NC) 6 , for example is made by passing chloride through 
a solution of potassium ferrocyanide, 

2K 4 Fe(NC) 6 -f C1 2 = 2KC1 + 2K 3 Fe(NC) 6 

It crystallizes in dark red prisms and is known as "red prussiate 
of potash." It is the potassium salt of the complex ferricyanogen 

ion, Fe(NC) 6 , which differs from the ferrocyanogen in that 

it carries one less charge. Ferric salts do not give any precipi- 
tate with ferricyanides, but the solution becomes somewhat dark 
brown. Ferrous salts, however, react and produce a dark blue 
precipitate of Turnbull's blue, ferrous ferricyanide, Fe 3 (Fe(NC) 6 )2, 
which is very similar to Prussian blue. In fact, some consider it 
identical with Prussian blue and think that ferricyanogen oxi- 
dizes the ferrous ion to the ferric- and is itself reduced to the 
ferrocyanogen before they combine. 

When solutions of potassium ferricyanide and ammonium 
ferric citrate are mixed, a yellowish-brown liquid results. If 
this be spread over paper and the latter dried in the dark, sen- 
sitized blue print paper results. When it is exposed to sunlight 
under a tracing or negative, it turns blue wherever the light 
strikes the paper, from the formation of insoluble blue compounds 
like Prussian blue. This may be due to the reduction of the 
ferric salt to the ferrous or of the ferricyanogen ion to the ferro- 
cyanogen and then their interaction. After the printing, the 
picture is "fixed" by washing in water which dissolves the 
unchanged mixture, leaving white lines on a blue ground. 

Oxalates. Ferrous oxalate, FeC 2 O4, is precipitated as a yel- 
low salt by the addition of oxalic acid to a ferrous solution. It 
is soluble in potassium oxalate owing to the formation of the 
complex salt, K^Fe^O^, which has a yellowish-red color not 
unlike that of many ferric salts. Ferric oxalate, unlike other 
ferric salts, has a green color remotely like that of the ferrous 
salts. It forms a stable complex salt with potassium oxalate, 
KsFeCC^O^s, which also has a green color and gives the com- 
plex ion Fe (204)3 . The usefulness of oxalic acid and of 

acid potassium oxalate in removing rust and ink spots from 
fabrics is due to the formation of this complex ion. Solutions 
of ferric oxalate and of its complex compounds are very sensitive 


to the light and. are quickly reduced to ferrous oxalate with the 
evolution of carbon dioxide. 

Ferrates. When ferric hydroxide is suspended in potassium 
hydroxide solution and chlorine passed in, a red solution of potas- 
sium ferrate, K2FeC>4, is formed. This may be obtained as 
crystals isomorphous with potassium sulfate and chromate. 
It is very unstable. Other salts have been prepared by 

The Corrosion of Iron. It is, of course, well known to every- 
one that iron when exposed to air and water will gradually be 
converted into partially dehydrated ferric hydroxide, which 
is called iron rust. The importance of an understanding of the 
conditions which influence this change is obvious, owing to the 
almost universal use of iron in our modern structures. There is 
a good deal of dispute concerning many of the facts in connection 
with this and still more as to their interpretation, but the fol- 
lowing observations are generally accepted as facts. First, iron 
will be rapidly dissolved and corroded by practically all dilute 
acids including carbonic, solutions of ferrous salts and hydrogen 
gas being the usual products. Second, iron will slowly dissolve 
in pure water even in the absence of oxygen or carbon dioxide. 
The action is not very extensive, but hydrogen gas is evolved and 
the solution becomes distinctly alkaline from the ferrous hy- 
droxide which is dissolved. In the entire absence of air, the solu- 
tion of ferrous hydroxide formed by the action of distilled water 
upon iron will remain clear; but if air be passed through it, the 
whole solution will take on the color of iron rust and will deposit 
it. Third, that the corrosion of iron will be greatly increased by 
the presence of oxygen. Fourth, that the iron is more readily 
attacked the more impure and non-homogeneous it is. Fifth, 
that in alkaline solutions the corrosion is greatly hindered or 
suppressed. Sixth, that in certain strong oxidizing agents, such 
as concentrated nitric acid, dichromates, and red lead, the iron 
becomes " passive" and will retain this condition for some time 
after the removal of the oxidizing agent. 

The most consistent interpretation of these facts seems to be 
the following: That the oxidation of the iron takes place in two 
stages; first, to the ferrous and then to the ferric condition; 
that the oxidizing agent which dissolves the iron to the ferrous 


state is hydrogen as ion either from acids or from water, and 
that the oxidation to the ferric is brought about by the oxygen of 
the air. Since an oxidizing agent is the more active the more 
concentrated it is, the more rapid solution in acids than in water 
is readily understood as being due to the smaller concentration 
of the ionized hydrogen in the latter. The protective action of 
alkaline solutions may be ascribed to the concentration of the 
hydrogen as ion having been so far reduced by the increase in the 
concentration of the hydroxyl that it is not a strong enough 
agent to oxidize the iron, and so the first stage of the rusting 
cannot take place. 

If the iron is non-homogeneous, voltaic couples will be set up at 
various points on the surface and local electrolytic action will 
result. Hydrogen gas will be deposited at those points where 
the metal for any cause is acting as cathode, and since this hydro- 
gen has a tendency to pass back into solution as ion, the local cells 
will be polarized and the action hindered. If oxygen is present, 
it will act as a depolarizer and hence increase the rate of electro- 
lytic corrosion, and then in addition will oxidize the ferrous ion 
so produced to the ferric, and so complete the rusting process. 
There seems to be no satisfactory explanation for the passivity 
of iron. 

Analytical Properties of Iron. Because of the precipitation 
of ferric hydroxide by ammonium hydroxide in the presence of 
ammonium salts, iron is placed in an analytical group with 
aluminum and chromium. Ferrous salts give a deep blue 
with ferricyanides, while ferric compounds give the same color 
with ferrocyanides and also a blood red with thiocyanates. With 
the borax bead, iron compounds give green colorations in the 
reducing flame and a faint yellow or brown in the oxidizing 


Cobalt has a slightly higher atomic weight than nickel, but 
it is more closely related to iron than nickel is, so it is usually 
placed ahead of the latter in the periodic system. Cobalt 
occurs in nature in rather small quantities chiefly as smaltite, 
CoAs 2 , and cobaltite, CoAsS, and is commonly associated 


with iron, nickel, and manganese. It is almost always a con- 
stituent of meteoric iron and is present in the sun as indicated 
by the spectrum. The metal may be obtained by the reduction 
of the oxide at a high temperature with carbon or hydrogen, 
by the Goldschmidt process, or by the electrolysis of a solu- 
tion of the sulfate. It is magnetic and has the color of polished 
iron, becoming pink on exposure to the air, and melts at 1,494. 
It slowly oxidizes when heated and burns at a high temperature. 
It is attacked by dilute acids, but tends to become passive as iron 
does. Very recent experiments indicate that cobalt can be sub- 
stituted with advantage for nickel in electroplating other metals. 
The coating may be applied much more rapidly, it is harder and 
more durable and hence may be made thinner than nickel with 
equally good results. Cobalt forms two series of compounds 
corresponding to ferrous and ferric. The simple cobaltous com- 
pounds are more stable than the simple cobaltic, but the case is 
the other way around with the complex compounds. 

Cobaltous Compounds. When sodium hydroxide is added 
to a cobaltous solution, a blue precipitate of basic salt is formed 
which on standing becomes pink and is changed to the hydroxide, 
Co(OH) 2 . This, like the corresponding ferrous, manganous, and 
magnesium compounds, is soluble in ammonium salts. In addi- 
tion, it is soluble in ammonium hydroxide presumably because 
of the formation of complex ammonia ions. The ammoniacal 
solution absorbs oxygen, and a series of complex cobaltic am- 
monia salts are formed which will be treated briefly a little later. 
The cobaltous salts in dilute solution have a red color and so do 
most of the fully hydrated solid salts, but the partially or com- 
pletely dehydrated compounds are generally blue. Cobalt 
chloride, CoCl 2 -6H 2 O, is especially sensitive in this respect. 
One may prepare a dilute solution of cobalt chloride and write 
with it upon paper; after drying, the delicate pink of the salt 
is scarcely visible, but when heated it loses water and the charac- 
ters become distinctly traced in the blue of the less hydrated 
salt. Upon standing in the air at ordinary temperatures, 
the salt takes up water once more and the writing disappears. 
This is the basis of the so-called "sympathetic ink." 

Cobalt Sulfate. Cobalt sulfate, CoS0 4 -7H 2 0, is isomorphous 
with magnesium sulfate and forms double sulfates with ammo- 


mum and potassium sulfates, K 2 Co(S04) 2 -6H 2 which are iso- 
morphous with the other double sulfates of the same type. 

Cobalt Nitrate. The nitrate, Co(NO 3 ) 2 -6H 2 O, is perhaps the 
most used cobalt salt, being employed in blowpipe analysis and 
in making other cobalt preparations. 

Cobalt Sulfide. Cobalt sulfide, CoS, is black and is precipi- 
tated by ammonium sulfide. It is peculiar in that it is not 
precipitated by hydrogen sulfide in acid solutions, except in the 
presence of a large amount of sodium acetate; but after the pre- 
cipitate has once been obtained it is scarcely appreciably dis- 
solved by cold dilute acids. It probably forms a much more 
stable modification soon after precipitation. 

Cobalt Glass. Compounds of cobalt color the borax bead a 
beautiful blue and impart to fused silicates the same color. 
Smalt is a kind of glass which is made by fusing together sand, 
crude cobalt oxides and potassium carbonate. When ground 
very fine it is used as a pigment and for coloring glass and porce- 
lain blue. When aluminum oxide or a salt of aluminum is 
heated with cobalt oxide or one of the cobalt salts, a blue com- 
pound is produced which is used as a paint. 

Cobaltic Compounds. Cobaltic hydroxide, Co(OH) 3 , is formed 
when a hypochlorite is added to a cobaltous salt. It is brownish- 
black and is soluble in cold hydrochloric acid, but the solution 
quickly decomposes with the evolution of chlorine and the for- 
mation of cobaltous chloride. The corresponding oxide, Co 2 O 3 , 
may be prepared by gently heating the nitrate; when more 
strongly heated it goes to the oxide, Co 3 O4. A cold acid solution 
of cobaltous sulfate may be oxidized at the anode of an elec- 
trolytic cell to cobaltic sulfate, Co 2 (S0 4 )3-18H 2 O.. Ammonium, 
potassium, rubidium or cesium sulfates form alums with thi s 

Complex Compounds. When potassium cyanide is added to 
a cobaltous solution, a precipitate of cobalt cyanide is formed, 
Co(NC) 2 , which dissolves in excess of potassium cyanide with the 
formation of potassium cobaltocyanide, K 4 Co(NC)e, analogous 
to potassium ferrocyanide. This is a very powerful reducing 
agent and in acid solutions will take up oxygen from the air and 
rapidly pass into potassium cobalticyanide K 3 Co(NC) 6 ; 
4K 4 Co(NC) 6 + 4HNC + O 2 = 4K 3 Co(NC) 6 + 4KNC + H 2 O 


The complex ion so formed is very stable and does not give the 
reactions of the Co +++ ion. The corresponding acid is also 
very stable. 

When acetic acid and a nitrite are added to a cobaltous salt, 
the latter is first oxidized to the cobaltic state and then the co- 
baltic ion combines with the nitrite to make the complex ion, 
Co(NO2)e This yields very difficultly soluble compounds 
with potassium and ammonium ions, and in the form of a solution 
of the sodium salt, it is much used as a reagent for potassium. 
Since nickel does not form any corresponding compound, this 
furnishes a means for the separation of the two elements. 

When cobaltous hydroxide is dissolved in ammonia and 
the solution exposed to the air or to the action of oxidizing agents, 
cobaltic ammonia compounds are formed. These act as though 
they were compounds of the complex cation Co(NH 3 ) n + + + , in 
which n is 3, 4, 5, or 6, with three atomic weights of a mono- 
valent anion so that the general formula of the compounds may 
be written, Co(NH 3 ) n A 3 , in which A represents any monovalent 
anion. The following rules will give a slight idea of the behavior 
of these compounds. When "n" is 6, all three atomic weights 
of A are easily and equally ionizable. When "n" is 5, two 
atomic weights of A are easily ionizable while the third is far less 
so. When "n" is 4 one atomic weight of A is easily ionizable 
and the other two are but slightly so. When "n" is 3 the 
substance is practically undissociated. These compounds may 
contain a mixture of anions in the same salt, and the ammonia 
may be replaced wholly or in part by water or a large number 
of organic compounds. These facts make the chemistry of these 
salts so complex that several hundred of the compounds are 


Nickel occurs with cobalt, iron, and manganese. A few ores 
are nicollite, NiAs, nickel glance, NiAsS, and genthite, Mg2Ni 2 - 
H 4 (Si04)3'4H 2 O. The chief source of the metal is Ontario, 
and New Caledonia. Nickel is more abundant than cobalt and 
hence is cheaper and has more applications. The commercial 
extraction of nickel from its ores is a very complex process and 


will not be described here, but something will be said of one 
method in connection with nickel carbonyl. 

Properties. Nickel is a silvery white metal with a tinge of 
yellow. It takes a high polish and is hard but malleable and 
may be rolled into thin sheets or drawn into fine wire. The pure 
metal melts at 1,451. It is magnetic. Nickel is a component 
of a number of important alloys. German silver contains ap- 
proximately 60 per cent, copper, 25 per cent, zinc, and 15 per 
cent, nickel, but the composition varies widely. Nickel coins 
usually contain 75 per cent, copper, and 25 per cent, nickel. 
Nickel steel has very valuable properties because of its increased 
strength and toughness. A steel containing about 36 per cent, of 
nickel has a very small temperature coefficient and is known as 
" invar." It is used in making standard meters because their 
length will vary but slightly with the temperature. A complex 
wire having a core of nickel steel, a jacket of copper and a coat 
of platinum over all has a slightly smaller coefficient of expansion 
than glass and is coming into use in place of the very precious 
metal platinum for sealing in wires in^electric light bulbs. Monel 
metal is a nickel-copper alloy obtained by the direct smelting of 
the nickel copper ores mined at Sudbury, Ontario. It can be 
cheaply produced and has the very valuable properties o